Process for sequestration of carbon dioxide and minerals from industrial waste products

Abstract

The present invention provides a process for producing calcium carbonate and magnesium hydroxide and / or magnesium carbonate, using an aqueous brine solution, comprising the steps of precipitating calcium carbonate and magnesium hydroxide and / or magnesium carbonate; and then removing the calcium carbonate and the magnesium hydroxide and / or the magnesium carbonate.

Description

FIELD OF THE INVENTION This invention relates to a process for sequestering industrially produced carbon dioxide and minerals from industrial waste products. In particular, dissolved carbon dioxide (mostly present as bicarbonate) and minerals may be extracted from waste brine produced by desalination plants and converted into industrially useful products. In addition, the process can operate in a way which incorporates carbon dioxide emitted from other industrial sources into industrially useful products. BACKGROUND OF THE INVENTION Carbon sequestration refers to the process of capturing and storing atmospheric or hydrospheric carbon dioxide which can be achieved through various chemical and biological processes, but also occurs naturally over long timescales. Industrial waste, such as desalination brine, contains reasonable quantities of metal ions. These ions also form carbonate minerals found in the Earth’s crust which are typically mined and processed at high temperatures and pressures into useful chemicals / compounds for industrial use. The cations within industrial waste sources have the capacity to sequester large quantities of carbon dioxide (as bicarbonates or carbonates) under ambient conditions and over comparatively short timescales, if the thermochemical properties (e.g. alkalinity, pH, temperature, and pressure) are controlled in situ to reduce the solubility of the target mineral. Authors have investigated the potential for the application of chemical precipitation techniques to desalination brine to mineralise carbon dioxide as magnesium-, and calciumbased carbonate minerals and to generate magnesium hydroxide. US8197696B1 discloses a method for separating magnesium hydroxide from calcium in hydroxide form or carbonate form or sulfate form from a saline stream. The method comprises a) converting the alkalinity of the saline stream to carbon dioxide by adding an acid to produce nearly alkalinity-free saline stream; (b) removing magnesium hydroxide from the nearly alkalinity-free saline stream; (c) removing calcium in the hydroxide form from the magnesium-free saline stream; or (d) removing calcium in the carbonate form from the magnesium-free saline stream; or (e) removing calcium in the sulfate form from the magnesium-free saline stream. US2014 / 263081A1 discloses various industrial embodiments for a desalination and water purification system. Specifically, the system includes a pre-treatment step that precipitates scale-forming compounds. In one embodiment, the method involves a first step of initial removal of magnesium ions by precipitating magnesium hydroxide at high pH, and then removing the precipitate by either sedimentation or filtering. Following magnesium hydroxide precipitation, carbonate ions are added in the form of CO2 sparging, to subsequently precipitate calcium, barium, and other divalent cations as carbonates by adjusting the pH to about 10.2 or greater. WO2010 / 118425A1 discloses a method for removing hydrocarbons and scale forming compounds from tap water, contaminated aqueous solutions, seawater, and saline brines, such as produce water, comprising the addition of carbonate ions by CO2 sparging, or divalent cations, so as precipitate calcium and magnesium carbonates by adjusting pH to about 10.2, thus permanently sequestering CO2 from the atmosphere, and then removing such precipitates sequentially for either sale of disposal. US2019 / 0232216A1 discloses a method of mineralisation of carbon dioxide. Specifically, the method comprises forming an alkaline in aqueous solution containing carbonate anions by dissolving carbon dioxide and an alkali such as ammonia in water. Next, the method comprises mixing the alkaline aqueous solution with a water source containing magnesium and calcium cations. A first product containing calcium cations and carbonate anions is precipitated in a first precipitation step at a first pH (e.g. around pH7.5) and then a second product containing magnesium cations and carbonate anions is precipitated in a second precipitation step at a second, higher pH e.g. around pH 9.5. US2011 / 0059000A1 discloses methods of sequestering carbon dioxide. Specifically it relates to treating at least a portion of flue gas emitted from a cement plant, wherein said flue gas comprises carbon dioxide, with a proton removing agent to produce a solution comprising carbonate ions; treating said solution with alkaline-earth-metal-ions to produce a precipitate comprising carbonate; and delivering at least a portion of said precipitate to said cement plant as a feedstock. US2010 / 0158786A1 discloses desalination methods that include carbonate compound precipitation. In particular, feed water and / or waste brine is subjected to carbonate compound precipitation conditions to produce a carbonate compound precipitate and an alkali-earthmetal-ion-depleted water. CN101157504A discloses a process for softening hardwater by employing SO2 and CO2 in exhaust gas. Specifically, the process includes allowing the SO2 and the CO2 contained in the exhaust gas to react with hard water; stopping the reaction when the pH of the mixed solution is tested to be 2.5-3.3; adjusting the pH of the mixed solution after reaction with alkaline solution to be 6.5-7.5; and performing filtration to obtain softened water and filter residue. EP140211 IB 1 discloses a method for treating process water circulating in a process for producing paper. Specifically, the method comprises separating a stream of process water from said aqueous cellulosic pulp suspension, said separated process water having an elevated level of free calcium ions, adding sodium hydroxide to said process water for providing therein, without the addition of calcium oxide or hydroxide, a pH of not less than pH 9, subsequently distributing gaseous carbon dioxide in said process water in an amount sufficient for lowering the pH of said water to a pH level of 8.5 to 6.0 for precipitating a substantial part of said calcium ions as solid calcium carbonate in said water and for obtaining a decreased level of free calcium ions in said water, optionally removing solid calcium carbonate from said process water. The resulting calcium ion depleted process water is then used in the papermaking process disclosed therein. US2021 / 0024364A1 discloses methods for producing hydrated calcium and magnesium salts from alkaline industrial wastes. Specifically, the divalent ion-containing solution is subjected to concentration to form a concentrated divalent ion-containing solution. Precipitation of a divalent ion hydroxide salt is induced from the concentrated divalent ioncontaining solution. In other cases, the concentrated divalent ion-containing solution is exposed to carbon dioxide to induce precipitation of a divalent ion carbonate salt. The present inventors previously found that earlier disclosed methods can be significantly improved by performing initial alkalization and nanofiltration steps before precipitation (see international application number PCT / GB2023 / 051727). This allowed for increasing the concentration of target minerals with low energy requirement to improve chemical precipitation rate, yield and reagent consumption. In particular, this process enabled the precipitation and isolation of calcium carbonate, and optionally magnesium hydroxide and / or magnesium carbonate. However, an unsolved problem in the art is maximising the yield and purity of precipitated calcium carbonate, magnesium hydroxide and magnesium carbonate from complex mixtures in aqueous brines. This is at least because it is exceptionally challenging to prevent co-precipitation when the brine is alkalized. The present invention solves this problem by performing a sequential precipitation of calcium carbonate, followed by the precipitation of magnesium hydroxide and / or magnesium carbonate. Specifically, the present inventors surprisingly found that controlling the order of precipitation with precise pH ranges provides the respective minerals in greater purity and yield than previous approaches, and also avoids the need for preceding processing steps such as nanofiltration. SUMMARY OF THE INVENTION In a first embodiment, the present invention provides a process for producing calcium carbonate and magnesium hydroxide, using an aqueous brine solution which contains calcium and magnesium, comprising the steps of: a. (I) adding an alkalinity source to the aqueous brine solution such that the pH of the solution is between 8.7 and 9.3; a. (II) adding carbon dioxide gas to the aqueous brine solution produced in step a(I) such that the pH of the solution is between 5.0 and 6.0; a. (Ill) adding an alkalinity source to the aqueous brine solution produced in step a(II) such that the pH of the solution is between 8.5 and 9.5 and calcium carbonate precipitates; b. removing calcium carbonate produced in step a(III); then c. (I) precipitating magnesium hydroxide by adding an alkalinity source to the calcium-depleted aqueous brine solution such that the pH of the solution is between 10.5 and 11.5 and magnesium hydroxide precipitates; and d. (I) removing magnesium hydroxide produced in step c(I). In a second embodiment, the invention provides a process for producing calcium carbonate and magnesium carbonate, using an aqueous brine solution which contains calcium and magnesium, comprising the steps of: a. (I) adding an alkalinity source to the aqueous brine solution such that the pH of the solution is between 8.7 and 9.3; a. (II) adding carbon dioxide gas to the aqueous brine solution produced in step a(I) such that the pH of the solution is between 5.0 and 6.0; a. (Ill) adding an alkalinity source to the aqueous brine solution produced in step a(II) such that the pH of the solution is between 8.5 and 9.5 and calcium carbonate precipitates; b. removing calcium carbonate produced in step a(III); then c. (II) precipitating magnesium carbonate by adding an alkalinity source and carbon dioxide gas to the calcium-depleted aqueous brine solution such that the pH of the solution is between 9.0 and 10 and magnesium carbonate precipitates; and d. (II) removing magnesium carbonate produced in step c(II). In a third embodiment, the invention provides a process for producing calcium carbonate, magnesium carbonate and magnesium hydroxide using an aqueous brine solution which contains calcium and magnesium, comprising the steps of: a. (I) adding an alkalinity source to the aqueous brine solution such that the pH of the solution is between 8.7 and 9.3; a. (II) adding carbon dioxide gas to the aqueous brine solution produced in step a(I) such that the pH of the solution is between 5.0 and 6.0; a. (Ill) adding an alkalinity source to the aqueous brine solution produced in step a(II) such that the pH of the solution is between 8.5 and 9.5 and calcium carbonate precipitates; b. removing calcium carbonate produced in step a(III); then c. (II) precipitating magnesium carbonate by adding an alkalinity source and carbon dioxide gas to the calcium-depleted aqueous brine solution such that the pH of the solution is between 9.0 and 10 and magnesium carbonate precipitates; and d. (II) removing magnesium carbonate produced in step c(II); then c. (I) precipitating magnesium hydroxide by adding an alkalinity source to the calcium / magnesium-depleted aqueous brine solution such that the pH of the solution is between 10.5 and 11.5 and magnesium hydroxide precipitates; and d. (I) removing magnesium hydroxide produced in step c(I). BRIEF DESCRIPTION OF THE DRAWINGS Examples of the invention will now be described in detail with reference to the accompanying drawings, in which: Figure 1 shows a schematic representation of the apparatus used for steps a(I)-a(III) and step b of Examples 1 and 2. Figure 2 shows a schematic representation of the apparatus used for step c(I) and d(I) of Example 1. Figure 3 shows a schematic representation of the apparatus used for step c(II) and d(II) of Example 2. Figure 4 shows the composition of the crude precipitate from steps a(I)-a(III) under conditions 1-4 of Example 1, measured by Inductively Coupled Plasma Optical Emission Spectroscopy (ICP-OES). Figure 5 shows the alkalinity of the brine from steps a(I)-a(III) under conditions 1-4 of Example 1. Figure 6 shows the composition of the solids from step b under conditions 1-4 of Example 1, measured by ICP-OES. Figure 7 shows the composition of the solids from step d(I) under conditions 1-5 of Example 1, measured by ICP-OES. Figure 8 shows the composition of the crude precipitate from step c(II) under conditions 1-4 of Example 2, measured by ICP-OES. Figure 9 shows the composition of solids from step d(II) under conditions 3-5 of Example 2, measured by ICP-OES. Figure 10a shows the total calcium and magnesium change between the column and overflow brine versus total NaOH consumption (as total NaOH kg) at pH 8.1, 8.5, 8.8, 9.1 for step a(III). Figure 10b shows the amount of solid filtered at pH 8.5, 8.8, 9.1 for step a(III). No sample was able to be filtered from pH 8.1 due to low precipitation Figure 11 shows the ICP-OES analysis of the solids composition obtained from precipitation at pH 8.5, 8.8 and 9.1 for step a(III). Figure 12 shows the concentration of magnesium and calcium in the brine at various locations during step a(I)-a(III) when using 10% NaOH (above) and l.lwt% lime (below). Figure 13 shows the ICP-OES analysis of the solids isolated from step a(I)-a(III) when using 10% NaOH pre-alkalinisation versus lime pre-alkalinisation. Figure 14 shows the change in magnesium concentration from step c(II) at pH precipitation setpoints, 8.75, 9.0 and 9.3. Figure 15 shows the yield and ICP-OES analysis of the solid isolated from 1 L of brine in step c(II) when varying the CO2 flow rate. Figure 16 shows the concentration of magnesium in the brine compared to overflow when performing step c(I)-d(I) after c(II)-d(II). Figure 17 shows the theoretical yield (based on Mg content in IL brine) compared to actual yield of magnesium hydroxide when performing step c(I)-d(I) after c(II)-d(II). Figure 18a shows the ICP-OES analysis of the filter press solids when performing step c(I)-d(I) after c(II)-d(II). Figure 18b shows the effect of repulping on the purity of magnesium hydroxide from c(I)-d(I) performed after c(II)-d(II). Purity determined by ICP-OES analysis. DETAILED DESCRIPTION OF THE EMBODIMENTS Step a(I): Alkalization The first step in the process of all three embodiments of the present invention involves increasing the pH of the aqueous brine solution to a value between 8.7 and 9.3. This converts bicarbonate ions into carbonate ions. HC03' + OH- CO32 + H2O Preferably, the pH increase in step a(I) is achieved by addition of an aqueous alkali reagent, such as sodium hydroxide, potassium hydroxide, calcium hydroxide (lime), magnesium hydroxide, sodium carbonate or ammonia, or alternatively, an alkali waste or waste water from an industrial process such as fly ash or lime sludge. In an alternative method the hydroxide ions can be produced through electrical methods, such as, electrolysis of water. Preferably the aqueous alkali reagent is an alkaline or alkali metal hydroxide, preferably sodium hydroxide or calcium hydroxide, most preferably 10 % w / w sodium hydroxide. It was found that care was required when increasing the pH, because if the pH in any localised part of the solution became too high then it caused precipitation. Thus, in a preferred embodiment, the alkalising agent is pre-mixed with a small quantity of the initial aqueous brine solution or other recycled process stream before being introduced into the main process stream. The alkalising source may also be a high pH waste stream from this process or another compatible industrial process. It was found that at pH of about 9, the calcium ions in the aqueous brine solution remained fully dissolved, but at 9.3 or above, there was a precipitation of the calcium (most likely as CaCO3). Ca2+ + CO32' CaCO3 (s) Thus, in a preferred embodiment, the pH of the aqueous brine solution is increased to a value from 8.8 to 9.2, still more preferably to a value from 8.9 to 9.1, most preferably to a value of about 9.0. Steps a(II) and a(III): Calcium carbonate precipitation The second and third steps in the process of all three embodiments of the present invention result in the precipitation of calcium carbonate which involves the reaction of calcium ions with carbonate ions. Ca2+ + CO32' CaCO3 It must be noted that the concentration of carbonate ions is also a function of pH with higher pH conditions converting bicarbonate ions into carbonate ions until a pH of 10.25 at which point most of the alkalinity is in carbonate form. H2CO3 H+ + HCO3- 2H+ + CO32' The carbonate present in the product of step a(I) is supplemented with carbonate generated by the reaction of CO2 (e.g. exhaust gas from industry) with the alkaline solution, thus forming a carbon capture, utilisation and storage system (CCUS). Initial experiments indicated that calcite precipitation could be optimised by adjustment of the pH. Thus, step a(III) involves adjusting the pH of the aqueous brine solution to a value between 8.5 and 9.5. Preferably the pH is adjusted to a value between 8.8 and 9.5. More preferably, the pH is adjusted to a value between 9.0 and 9.5. More preferably, the pH is adjusted to a value between 9.0 and 9.2. Most preferably, the pH is adjusted to a value of about 9.1. Preferably, sodium hydroxide is used as the alkalinity source. When CO2 is used as the source of carbonate (i.e. step a(II)) it may involve bubbling CO2 through the aqueous brine solution to provide a solution with a pH between 5.0 and 6.0, more preferably between 5.2 and 5.6, most preferably to about 5.4. It may involve bubbling CO2 directly into the aqueous brine solution, at a flow rate of between 6 and 14 litres per minute, preferably 7.5 litres per minute. This is equivalent to a flow rate of between 0.54 and 1.00 L CO2 / L Brine, preferably 0.68 L CO2 / L Brine. Step a(III) involves adjusting the pH of the aqueous brine solution. This is preferably carried out by adding a hydroxide solution (e.g. sodium hydroxide) having a concentration in the range from 0.1 to 2.0 M. More preferably, this is carried out by adding a hydroxide solution (e.g. sodium hydroxide) having a concentration in the range from 0.1 to 1.5 M. Most preferably, this is carried out by adding a hydroxide solution (e.g. sodium hydroxide) having a concentration of about 0.15 M. Initial experiments also indicated that calcite purity could be improved by seeded precipitation. Thus, in a preferred embodiment, prior to or during step a(III), calcium carbonate is added to the aqueous brine solution in an amount sufficient to promote precipitation of calcium carbonate. The seed can be either a ground calcium carbonate product added separately, or a precipitated calcium carbonate either isolated or recirculated from this process. Preferably, seed calcite is added in an amount from 0.1 to 20 % by mass based on the theoretical yield of calcium carbonate, more preferably from 10 to 20 % and most preferably about 15 %. Use of about 15 % seed calcite resulted in residual calcium levels which were low (around the 30-100 mg / L range) meaning that most of the calcium precipitated out as a solid. The yield, purity, and filterability were all good and improved relative to the unseeded experiments. Combining the preferred embodiments described above enabled calcite precipitation with greater than 90% recovery and purity (see Figures 4 and 6). In the context of the present invention, the term “calcium carbonate” encompasses amorphous and crystalline forms of anhydrous and hydrated forms of calcium carbonate. For example, the calcium carbonate may be an amorphous or crystalline form of one or more of CaCOs (e.g. calcite, aragonite or vaterite), CaCO3.H2O (e.g. monohydrocalcite), and CaCO3.6H2O (e.g. ikaite). Step b: Calcium carbonate removal The fourth step in the process of all three embodiments of the present invention is the removal of calcium carbonate from the system. This step may be carried out by any suitable method for removing solids from liquids. Preferably, the concentration of bicarbonate (HCO3 ) ions in the aqueous solution remaining after step b is 200 mg / L or less and / or the concentration of carbonate (CO32) ions in the aqueous solution remaining after step b is 200 mg / L or less. In one embodiment of the process, steps a(III) and b are repeated until the concentrations of bicarbonate (HCO3) and carbonate (CO32) ions in the aqueous solution remaining after step b are reduced to the desired levels. Filterability was found to be an issue due to particle size and wetness of product. To improve this, a flocculant may be used. A preferred flocculant is FLOP AM™ AN913 PWG which is commercially available from the company SNF FA (France). Preferably the removal of calcium carbonate comprises the steps of: (i) allowing the calcium carbonate and / or magnesium hydroxide or magnesium carbonate precipitate to settle; (ii) removing the supernatant liquid; and (iii) filtering the remaining slurry. Step c(I): Magnesium hydroxide precipitation The fifth step in the process of the first embodiment of the present invention, and the eighth step in the process of the third embodiment of the present invention, is the precipitation of magnesium hydroxide which involves the reaction of magnesium ions with hydroxide ions. Mg2+ + 20H-  Mg(0H)2 Precipitation of magnesium hydroxide is achieved by adding an alkalinity source, preferably sodium hydroxide, to the aqueous brine solution such that the pH of the solution is between 10.5 and 11.5, preferably between 10.8 and 11.2, more preferably between 10.8 and 11.0, most preferably to about 11.0, and magnesium hydroxide precipitates. Preferably, the amount of magnesium remaining in the aqueous brine solution after step c(I) is less than 200 ppm. In the context of the present invention, the term “magnesium hydroxide” encompasses amorphous and crystalline forms of anhydrous and hydrated forms of magnesium hydroxide. For example, the magnesium hydroxide may be an amorphous or crystalline form of Mg(0H)2 (e.g. brucite). Step d(I): Magnesium hydroxide removal The sixth step in the process of the first embodiment of the present invention, and the ninth step in the process of the third embodiment of the present invention, is the removal of magnesium hydroxide from the system. This step may be carried out by any suitable method for removing solids from liquids. To improve filterability, a flocculant may be used. A preferred flocculant is FLOP AM™ AN913 PWG which is commercially available from the company SNF FA (France). Preferably the removal of magnesium hydroxide comprises the steps of (i) allowing the magnesium hydroxide precipitate to settle; (ii) removing the supernatant liquid; and (iii) filtering the remaining slurry. Step c(II): Magnesium carbonate precipitation The fifth step in the process of the second and third embodiments of the present invention is the precipitation of magnesium carbonate which involves the reaction of magnesium ions with carbonate ions. Mg2+ + CO32' MgCO3 Precipitation of magnesium carbonate is achieved by adding an alkalinity source, preferably sodium hydroxide, and carbon dioxide gas to the aqueous brine solution such that the pH of the solution is between 9.0 and 10.0, preferably between 9.3 and 9.7, most preferably between 9.3 and 9.5, for example to about 9.3, 9.4 or 9.5, and magnesium carbonate precipitates. Preferably, the amount of magnesium remaining in the aqueous brine solution after step c(II) is less than 200 ppm. Preferably the carbon dioxide gas and alkalinity source are added simultaneously to the calcium-depleted aqueous brine solution produced in step a such that the pH of the solution is between 9.0 and 10.0 and magnesium carbonate precipitates. Preferably, adding carbon dioxide in step c(II) is at a flow rate of 2 L CO2 / L Brine. In the context of the present invention, the term “magnesium carbonate” encompasses amorphous and crystalline forms of anhydrous and hydrated forms of magnesium carbonate. For example, the magnesium carbonate may be an amorphous or crystalline form of one or more of MgCO3 (e.g. magnesite), MgCO3.2H2O (e.g. barringtonite), MgCO3.3H2O (e.g. nesquehonite), Mgs(CO3)4(OH)2 AH2O (e.g. hydromagnesite), Mg5(CO3)4(OH)2 5H2O (dypingite) and MgCO3.5H2O (e.g. lansfordite). Step d(II): Magnesium carbonate removal The sixth step in the process of the second and third embodiment of the present invention is the removal of magnesium carbonate from the system. This step may be carried out by any suitable method for removing solids from liquids. To improve filterability, a flocculant may be used. A preferred flocculant is FLOP AM™ AN913 PWG which is commercially available from the company SNF FA (France). Preferably the removal of magnesium carbonate comprises the steps of: (i) allowing the magnesium carbonate precipitate to settle; (ii) removing the supernatant liquid; and (iii) filtering the remaining slurry. EXAMPLES (GENERAL METHODOLOGY) Step a(T): Pre-alkalization Materials: Feed water was desalination brine sourced from the Adelaide Desalination Plant (operated and maintained by Adelaide Aqua Pty Ltd which is a subsidiary JV of Acciona in South Australia). 10% aqueous sodium hydroxide solution was sourced from ACE Chemicals. Experimental: Brine (e.g. 3500-5000 L / day) is fed into a day tank (such as a Bushmans Im3 HDPE feed water tank, fitted with an Intermediate Bulk Container (IBC) mixer and pH monitor). The pH of the brine is then adjusted to within the range of 8.5 to 9 by feeding in an alkalization agent (e.g. 10% NaOH) from an alkali tank (e.g. an IBC lOOOL HDPE tank). The flow rate of the alkalization agent is controlled by a dosing pump (e.g. Blackstone 316SS, model BL7016) in combination with a controller (e.g. Hanna instruments, model B79-16). Steps a(II) and a(III): Calcium carbonate precipitation (Figure 1) Materials: Carbon dioxide gas was sourced from BOC Limited. 25% and 50% aqueous sodium hydroxide solution were sourced from ACE Chemicals. Experimental: The alkalized brine from step a(I) is fed at a rate of 10-11 litres per minute (LPM) through a carbon dioxide adsorption column, with carbon dioxide gas bubbled through at a flow rate of about 7.5 LPM to provide a carbonated solution with a pH in the range of 5.2-5.6. This is achieved in combination with a CO2 pH monitor and controller. The resulting carbonated solution then flows to a precipitation tank (e.g. Rotamould ISOOL HDPE tank) at a rate of 10-11 LPM, wherein it is combined with a source of alkalinity (such as 25% or 50% NaOH) at a rate to achieve a pH of 8.5 to 9.5 (for example 9.0 to 9.4). The flow rate of the alkali source is remotely controlled with a controller (e.g. Walchem WPHPW 120 HA-N), in combination with a pH monitor (e.g. CWC M-10-A-05M). The resulting precipitate then overflows into a settling tank (e.g. Rotamould 2400L HDPE tank, fitted with a Phathom TSS probe S20-SWW-880-PP-10-MB) and is optionally treated with a flocculant (e.g. FLOP AM™ AN913 PWG manufactured by SNF FA (France)), concentration 0.025 g / L at a dose rate in the range of 2 to 4 L / h) to thicken and settle the precipitated solids. After a settling time of about 1 minute, the layers are separated, and the calcium carbonate slurry continues to step b. The supernatant is retained for step c(I) or c(II). Step b: Isolation of calcium carbonate (Figure 1) Experimental: The calcium carbonate slurry collects in a filter feed tank (e.g. Polymaster HDPE 5 4000 L tank). Optionally, the residual supernatant is separated from the precipitate using a Lamella thickener, or a rotary or drum thickener (e.g. Leopold Texler™ Lamella Clarifier). The slurry subsequently passes to a filter press (e.g. MBA470 manufactured by Innovative Filtration Solutions Pty Ltd) fitted with an air compressor (e.g. IRONAIR 3HP 100L compressor), whereby the calcium carbonate is filtered and washed with tap water (-100 L) at 2 bar, before increasing the pressure to 7 bar to remove the excess water. The pressure is then maintained, and the calcium carbonate is rinsed and dried two further times (-200 L) to provide a solid with a moisture content of about 35-45%. Step c(D: Magnesium hydroxide precipitation (Figure 2) Experimental: The supernatant from step a flows to the precipitation tank (e.g. Rotamould 1500L HDPE tank) at a rate of 10-11 LPM, wherein it is combined with a source of alkalinity (such as 25% or 50 % NaOH) to achieve a pH in the range of 10.5-11.5. The flow rates of the alkali source and the supernatant are remotely controlled with a controller (e.g. Walchem WPHPW 120 HA-N), in combination with a pH monitor (e.g. CWC M-10-A-05M). The resulting precipitate then overflows into a settling tank (e.g. Rotamould 2400L HDPE tank, fitted with a Phathom TSS probe S20-SWW-880-PP-10-MB) and is optionally treated with a flocculant (e.g. FLOP AM™ AN913 PWG manufactured by SNF FA (France)), concentration 0.025 g / L at a dose rate in the range of 2 to 4 L / h to thicken and settle the precipitated solids After a settling time of about 1 minute, the layers are separated, and the magnesium hydroxide slurry continues to step d(I). The supernatant is then directed to waste. Step d(I): Isolation of magnesium hydroxide (Figure 2) Experimental: The magnesium hydroxide slurry collects in a filter feed tank (e.g. Polymaster HDPE 5 4000 L tank). Optionally, the residual supernatant is separated from the precipitate using a Lamella thickener, or a rotary or drum thickener (e.g. Leopold Texler™ Lamella Clarifier). The slurry subsequently passes to a filter press (e.g. MBA470 manufactured by Innovative Filtration Solutions Pty Ltd) fitted with an air compressor (e.g. IRONAIR 3HP 100L compressor), whereby the magnesium hydroxide is filtered and washed with tap water (-100 L) at 2 bar, before increasing the pressure to 7 bar to remove the excess water. The pressure is then maintained, and the magnesium hydroxide is rinsed and dried two further times (-200 L) to provide a solid with a moisture content of 50-75% (e.g. 66%). Step c(II): Magnesium carbonate precipitation (Figure 3) Experimental: The supernatant from step a flows to the precipitation tank (e.g. Rotamould 1500L HDPE tank) at a rate of about 8-9 LPM. Carbon dioxide (sourced from BOC Limited) is bubbled through a diffuser at the bottom of the tank at a rate of 20 LPM, and an alkali source (such as 25% or 50% NaOH) is added through the top of the tank to maintain a pH in the range of 9-10. This is achieved with a controller (e.g. Walchem WPHPW 120 HA-N), in combination with a pH monitor (e.g. CWC M-10-A-05M). The resulting precipitate then overflows into a settling tank (e.g. Rotamould 2400L HDPE tank, fitted with a Phathom TSS probe S20-SWW-880-PP-10-MB and is optionally treated with a flocculant (e.g. FLOP AM™ AN913 PWG manufactured by SNF FA (France)), concentration 0.025 g / L at a dose rate in the range of 2 to 4 L / h to thicken and settle the precipitated solids After a settling time of about 1 minute, the layers are separated, and the magnesium carbonate slurry continues to step d(II). The supernatant is then directed to waste. Step d(II): Isolation of Magnesium carbonate (Figure 3) Experimental: The magnesium carbonate slurry collects in a filter feed tank (e.g. Polymaster HDPE 5 4000 L tank). Optionally, the residual supernatant is separated from the precipitate using a Lamella thickener, or a rotary or drum thickener (e.g. Leopold Texler™ Lamella Clarifier). The slurry subsequently passes to a filter press (e.g. MBA470 manufactured by Innovative Filtration Solutions Pty Ltd) fitted with an air compressor (e.g. IRONAIR 3HP 100L compressor), whereby the magnesium carbonate is filtered and washed with tap water (-100 L) at 2 bar, before increasing the pressure to 7 bar to remove the excess water. The pressure is then maintained, and the magnesium hydroxide is rinsed and dried two further times (-200 L) to provide a solid with a moisture content of 55-65%. EXAMPLES (SPECIFIC EXPERIMENTS) Example 1 In a first example, a process according to the general methodology outlined above was conducted in a pilot plant to investigate its suitability for medium to large scale sequestration of carbon dioxide and minerals from industrial waste products to obtain calcium carbonate and magnesium hydroxide. Starting with desalination brine sourced from the Adelaide Desalination Plant (40005000 L), the sequence of steps and apparatus used were steps a(I), a(II), a(III), b, c(I) and d(I). Steps a(I), a(II) and a(III) In accordance with the general methodology for step a(I) above, brine (4000-5000 L) was treated with 10% NaOH to afford an alkalized brine with a pH of 9. In accordance with the general methodology for steps a(II) and a(III) above, the alkalized brine from step a(I) (4000-5000 L) was fed at a rate of 10 LPM through the carbon dioxide adsorption column with carbon dioxide being bubbled through at a constant rate of 7.5 LPM or 12.5 LPM. This dropped the pH to 5.2-5.6. The carbonated solution then flowed to the precipitation tank to which 50% NaOH was added until a pH of 9.1, 9.3 or 9.5 was reached. This caused a precipitate to form which then overflowed into the settling tank with the addition of FLOP AM™ AN 913 PWG (0.025 g / L at a dose rate of 2 L / h). Four combinations of conditions (1-4) were tested as follows: Condition: pH Flow rate of CO2 (LPM) 1 9.5 7.5 2 9.3 7.5 3 9.1 7.5 4 9.1 12.5 The precipitates for each of the above conditions were left to settle from the supernatant, and then the precipitate slurries were transferred to the filter feed tank for step b. The supernatants were retained for step c(I). Samples of the precipitates (IL) from the precipitation tank were removed and filtered, rinsing with deionised water (3 x 100 mL). 0.4 g of the resulting solid was weighed, dissolved in 10% nitric acid (10-20 mL) and diluted with deionised water up to a total volume of 500 mL. The solutions were then analysed by ICP-OES. Composition of crude precipitate from conditions 1-4 (Figure 4): Condition: Calcium (%) Magnesium (%) Sodium (%) 1 (pH 9.5, 7.5 LPM) 95 4 1 2 (pH 9.3, 7.5 LPM) 93 4 3 3 (pH 9.1, 7.5 LPM) 97 1 1 4 (pH 9.1, 12.5 LPM) 98 1 1 ICP-OES revealed that the crude precipitates had the compositions outlined above (see also Figure 4). Notably, all pH levels provided a precipitate that was primarily calcium (93% or more) with very little magnesium (4% or less). Alkalinity of the brine from conditions 1-4 (Figure 5): Condition: Phenolphthalein alkalinity (mg / L) Total alkalinity (mg / L) 1 (pH 9.5, 7.5 LPM) 879 1658 2 (pH 9.3, 7.5 LPM) 883 1861 3 (pH 9.1, 7.5 LPM) 791 1760 4 (pH 9.1, 12.5 LPM) 912 2120 Furthermore, measurement of the alkalinity (see also Figure 5) of the brine indicatec that the amount of CO32' was higher at pH >9.3 (conditions 1 and 2) therefore these had a higher pH in the overflow. In contrast, pH 9.1 (conditions 3 and 4) had high total alkalinity but lower CO32 / HCO3' ratio, indicating that the alkalinity was largely made up of HCO3'. This held true even with the higher CO2 flow rate (condition 4). Overall, this indicated that lowering the pH provided greater purity of the CaCO3 in the slurry, and therefore pH 9.1 was most preferred. Step b In accordance with the general methodology for step b above, the precipitate slurries from each run (conditions 1-4) were transferred from the filter feed tank into the filter press until a pressure above 3 bar was reached and held for 1 minute. The solids obtained were analysed by ICP-OES and their compositions were as follows: Composition of solids from conditions 1-4 (Figure 6): Condition CaCO3 (%) Mg(OH)2(%) NaCl (%) NaSO-i(%) KC1 (%) 1 89.28 8.42 1.96 0.28 0.06 2 94.68 4.33 0.89 0.07 0.03 3 and 4 97.33 1.62 0.99 0.04 0.03 ICP-OES analysis (see also Figure 6) revealed that under all the above conditions, the composition of the solid was primarily calcium carbonate (>89%) with very little magnesium hydroxide (<9%). The highest degree of purity was observed at pH 9.1, with the solid obtained being >97% calcium carbonate. After drying the solid at 110 °C, XRD analysis confirmed that the isolated calcium carbonate was monohydrocalcite (CaCO3.H2O). Taken together, the ICP- OES analysis of the supernatant and solid confirmed that the most preferred conditions for step a(III) was a pH of about 9.1. Step c(l) (Figure 2) In accordance with the general methodology for step c(I) above, the supernatant from step a (4500-5000 L) was fed at a rate of about 10 LPM to the precipitation tank to which 50% NaOH was added until a pH of 9.75, 10.0, 10.25, 10.5 or 11.0 was reached. This caused a precipitate to form which then overflowed into the settling tank with the addition of FLOP AM™ AN 913 PWG flocculant (0.025 g / L at a dose rate of 3.5 L / h). The precipitate was left to settle from the supernatant, and then the precipitate slurry was transferred to the filter feed tank for step d(I). The five conditions (1-5) tested that varied the pH for precipitation were as follows: Condition: pH Magnesium precipitated (%) 1 9.75 3.5 2 10.0 32.5 3 10.25 79.1 4 10.5 95 5 11.0 99.7 The proportion of magnesium precipitated was calculated from the change in hardness in the brine before and after step c(I) (measured by titration with EDTA). This revealed that the greatest recovery of magnesium occurred at pH 11, the most preferred condition. Step d(l) (Figure 2) In accordance with the general methodology for step d(I) above, the precipitate slurries from each run (conditions 1-5) from step c(I) were transferred from the filter feed tank into the filter press until a pressure above 3 bar was reached and held for 1 minute. The solids obtained were analysed by ICP-OES and their composition were as follows: Composition of solids from conditions 1-5 (Figure 7): Condition CaCOs (%) Mg(OH)2(%) NaCl (%) KC1 (%) 1 10.97 71.51 15.57 1.49 0.46 2 1.58 76.56 19.99 1.29 0.59 3 3.96 74.90 18.16 2.44 0.55 4 1.73 80.13 15.94 1.68 0.53 5 1.81 98.04 0.10 0.04 0.00 ICP-OES analysis (see also Figure 7) revealed that unc er all the above conditions, the composition of the solid was primarily magnesium hydroxide (>71%) with very little calcium carbonate (<11%). Notably, the highest degree of purity was observed at pH 11, with the solid obtained being >98% magnesium hydroxide. XRD analysis confirmed that at pH 11, the isolated magnesium hydroxide was brucite with a purity of 98.25% (with 1.75% aragonite). Taken together, the most preferred condition for step c(I) was a pH of about 11, which provided the greatest recovery and purity of magnesium hydroxide. Example 2 In a second example, a process according to the embodiments and steps outlined above was conducted in a pilot plant to investigate its suitability for medium to large scale sequestration of carbon dioxide and minerals from industrial waste products to obtain calcium carbonate and magnesium carbonate. Starting with desalination brine sourced from the Adelaide Desalination Plant (35004000 L), the sequence of steps and apparatus used were steps a(I), a(II), a(III), b, c(II) and d(II). Steps a(I), a(II), a(III) and b (Figure 1) Steps a(I), a(II), a(III) and b were conducted identically to Example 1, wherein 3500-4000L of brine was used. Step c(ll) (Figure 3) In accordance with the general methodology for step c(II) above, the supernatant from step a (3500-4000 L) was fed at a rate of 8-9 LPM to the precipitation tank, to which CO2 was added at a rate of 20 LPM, as well as 50% NaOH at a rate to maintain a pH of 9.25, 9.5, 9.75, or 10.0. This caused a precipitate to form which then overflowed into the settling tank with the addition of FLOP AM™ AN 913 PWG flocculant (0.025 g / L at a dose rate of 4 L / h). The four conditions (1-4) tested that varied the pH for precipitation were as follows: Condition: 1 9.25 2 9.5 3 9.75 4 10.0 The precipitate was left to settle from the supernatant, and then the precipitate slurry was transferred to the filter feed tank for step d(II). Samples of the precipitates (IL) from the precipitation tank were removed and filtered, rinsing with deionised water (3 x 100 mL), and dried at 50 °C for 16 hours. 1 g of the resulting solid was weighed, dissolved in 10% nitric acid (10-20 mL) and diluted with deionised water up to a total volume of 500 mL. The solutions were then analysed by ICP-OES. Composition of crude precipitate from conditions 1-4 (Figure 8): Condition: Potassium (%) Sodium (%) Magnesium (%) Calcium (%) Chloride (%) SOT (%) 1 (pH 9.25) 0 0.29 92.31 7.40 0 0 2 (pH 9.5) 0 0.39 96.30 3.32 0 0 3 (pH 9.75) 0 4.18 92.04 0.68 3.10 0 4 (pH 10) 0.17 11.86 77.40 2.89 7.11 0.57 ICP-OES revealed (see also Figure 8) that the crude precipitates had the compositions outlined above. Notably, all pH levels provided a precipitate that was primarily magnesium (greater than 77 %) with very little calcium (less than 8%). The most preferred condition for step c(II) was a pH of about 9.5, as it provided the highest crude purity solid magnesium carbonate. Step d(ll) (Figure 3) In accordance with the general methodology for step d(II) above, the precipitate slurries from conditions 3 and 4 (pH 9.75 and 10) from step c(II) were transferred from the filter feed tank into the filter press until a pressure above 4 bar was reached and held for 1 minute. In a further experiment, condition 3 was repeated (conditions 5 herein) in accordance with the general methodology for step d(II) above, wherein the slurry was allowed to settle in the filter feed tank, and the residual supernatant was then removed using a drum pump before transferring to the filter press until a pressure above 5 bar was reached and held for 1 minute. The solids obtained were analysed by ICP-OES and their composition were as follows: Compositions of solidfrom conditions 3-5 (Figure 9): Condition: Potassium (%) Sodium (%) Magnesium (%) Calcium (%) Chloride (%) SOT (%) 3 0.27 12.32 76.55 0.88 9.54 0.44 4 0.49 23.17 53.71 4.89 16.87 0.89 5 0 0.15 99.01 0.76 0.08 0 ICP-OES analysis (see also Figure 9) revealed that conditions 3 and 4 provided solids that were primarily magnesium carbonate (>76% and >53%, respectively) with very little calcium carbonate (<1% and <5%, respectively). Notably, removing residual supernatant prior to transferring to the filter press resulted in a significantly improved purity of the magnesium carbonate, which was >99% (condition 5). XRD analysis confirmed that the isolated magnesium carbonate was the monohydrate (MgCO3.H2O). Taken together, this experiment demonstrated that optionally separating residual supernatant from the solid in the slurry prior to transferring to the filter press (i.e. by Lamella thickener or a rotary or drum thickener) can further improve the purity of the final solid. Example 3 In a third example, specific parameters of the process outlined in the embodiments above and in Examples 1 and 2 were further investigated. Investigating lower pH values for calcium carbonate precipitation - step a(III) As discussed in Example 1 above, pH values 9.5, 9.3 and 9.1 were trialled, and it was concluded that a pH of 9.1 was preferred as it provided greater purity of the CaCOs in the slurry. Therefore, in a further experiment, the pH for CaCOs precipitation was investigated at pH values of 8.1, 8.5, 8.8 and 9.1, to determine a preferred lower limit of pH needed to precipitate the maximum CaCOs, with the lowest consumption of alkalinity source. In accordance with the general procedure for step a(III) above (with 25% NaOH as the alkalinity source), pH values of 8.1, 8.5, 8.8 and 9.1 were trialled, with the change in calcium and magnesium concentration, and mass filtered, reported below in Figures 10A and 10B: Total Calcium and Magnesium change between the column and overflow brine compared with the total NaOH consumption at each pH setpoint (Figures 10A and 10B). Condition Change in brine concentration of Calcium (%) Change in brine concentration of Magnesium (%) Mass Filtered (g / L) NaOH consumption (kg / nC) 1 (pH=9.1) 91 1 1.666 0.86 2 (pH=8.8) 91 2 1.147 0.79 3 (pH=8.5) 76 0 0.761 0.77 4 (pH=8.1) 58 3 n.d. 0.88 When the pH set point was 8.1, minimal visible precipitation was observed in the tank, with only 58% of the calcium precipitating. At this low level of precipitation, filtration of a solid sample was not possible. Further precipitation was observed at a pH set point of 8.5, but the greatest increase was observed at the pH setpoint of 8.8, upwards of -90%. The mass filtered at pH 8.5 from IL of brine equated to 0.76 g / L. Notably, the percentage of calcium precipitated at pH 8.8 and 9.1 was the same, but pH 9.1 consumed slightly more NaOH per cubic meter of brine processed. The largest amount of solid was collected at pH 9.1. The composition of solid precipitates isolated at pH 8.5, 8.8 and 9.1 was then determined by ICP-OES twice, with the results reported in Figure 11 below. ICP-OES analysis of solids composition obtained from precipitation at pH 8.5, 8.8 and 9.1 (Figure 11) PH Ca (%) Mg (%) Na (%) K(%) Cl(%) SO / - (%) 9.1 sample 1 97 1 1 0 1 1 9.1 sample 2 97 2 1 0 0 0 8.8 sample 1 96 2 2 0 0 0 8.8 sample 2 96 2 2 0 0 0 8.5 sample 1 94 3 3 0 0 0 8.5 sample 2 91 3 5 0 0 0 The results show that for both samples collected at pH 8.8 and 9.1, the solids had high Ca content indicating a high purity of CaCOs, whereas slightly lower purity was observed at pH 8.5. Taken together, the results indicated that a pH of 8.8 for calcium precipitation was a preferred lower limit, as it still provided a good yield of calcium carbonate, with high purity, but lower NaOH consumption than at higher pH values. Investigatingpre-alkalinisation with lime as an alternative to sodium hydroxide - step a(I) In accordance with the general methodology for step a(I) above, lime was investigated as a replacement to 10% NaOH which was previously used for the pre-alkalinisation step. The solution of 1.1 wt% lime was made from combining powdered Ca(OH)2 and tap water. For this comparative experiment, the set pH for the day tank was 9 for step a(I). The subsequently alkalinised brines were then subjected to the general methodology for steps a(II) and a(III) above, whereby the carbon dioxide was bubbled through at a constant rate of 7.5 LPM, and the precipitation of calcium carbonate was performed at pH 9.1 with 25% NaOH. Comparison of the brine composition when using lime (condition 1) or sodium hydroxide (condition 2) in the pre-alkalinisation step is reported below in Figure 12. Brine analysis during each stage of steps a(I)-a(III) when using lime or sodium hydroxide for pre-alkalinisation (Figure 12) Condition for pre-alkalinisation (step a(I)) 1.1 wt% lime (condition 1) 10% NaOH (condition 2) Concentration of mineral Calcium (mg / L) Magnesium (mg / L) Calcium (mg / L) Magnesium (mg / L) Starting Brine 730.5 2280 733 2170 Day tank (step a(I)) 677.5 2210 704 2140 Column (step a(II)) 774.5 2265 731 2190 Overflow (step a(III)) 44.5 2295 75 2170 It can be seen above that composition of the brine did not substantially differ when using lime or sodium hydroxide for pre-alkalinisation. The concentration of magnesium was largely unchanged for both experiments, which is consistent with steps a(I)-(III) being highly selective for calcium carbonate precipitation. The overflow concentration of calcium indicated that most of the calcium was precipitated in both experiments. The consumption of base across steps a(I)-(III) is reported below. Base consumption in the precipitation of CaCOs using lime versus sodium hydroxide Step Base Base Consumption (kg / m ’) a(I) (condition 1) 1.1% lime 0.13 a(II)-a(III) (condition 1) 25% NaOH 0.99 a(I) (condition 2) 10% NaOH 0.06 a(II)-a(III) (condition 2) 25% NaOH 0.86 Following the precipitation step, the isolated solids were analysed by ICP-OES, which is reported in Figure 13 below. ICP-OES analysis of solids (Figure 13) Condition CaCOs (%) Mg(OH)2(%) NaCl (%) NaSO4 (%) KC1 (%) 1 (NaOH) 97.8 1.4 0.7 0.1 0.1 2 (lime) 99 0.8 0.1 0 0 Notably, the composition of the filter press solids were similar, with a respective purity of 97.8% and 99.0% CaCOs (Figure 13). Taken together, this experimented that lime as a base is a suitable alternative to sodium hydroxide for the pre-alkalinisation step a(I). Investigating lower pH values for magnesium carbonate precipitation - step c(II) As discussed in Example 2 above, pH values of 9.25, 9.5, 9.75 and 10 were trialled for magnesium carbonate precipitation (step c(II)), and it was concluded that a pH of about 9.5 was a preferred value, as it provided the highest crude purity solid magnesium carbonate. In said example, it was noted that at pH 9.25, magnesium carbonate could be precipitated using significantly less NaOH (3.2 kg / m3) compared to pH 9.5 (5.2 kg / m3), while having similar yields of precipitated magnesium (71-74%). Therefore, in a further experiment, lower pH values for magnesium carbonate precipitation were investigated. In accordance with the general methodology for step c(II) above, the calcium depleted brine from step a(I)-(III) was fed at a rate of 9-10 LPM to the precipitation tank, to which CO2 was added at a rate of 20 LPM, as well as 25% NaOH at a rate to maintain a pH of 8.75 (condition 1), 9.0 (condition 2) or 9.3 (condition 3). The measured change in magnesium concentration is reported below in Figure 14: Change in magnesium concentration at pH setpoints 8.75, 9.0 and 9.3 (Figure 14) Condition Magnesium decrease (%) 1 (pH 8.75) 9 (n=l) 2 (pH 9.0) 5 (mean, n=2) 3 (pH 9.3) 54 (mean, n=2) When analysing the change in magnesium concentration, at pH less than 9, less than 10% of magnesium was precipitated, whereas at pH 9.3, the precipitation was greater than 50% magnesium (Figure 14). Qualitatively speaking, at lower pH points of 8.75 and 9, the solution following addition of base was not very cloudy, visually indicating a low precipitation magnesium salts. This indicated that a particularly preferred pH for magnesium carbonate precipitation is 9.3. Investigating the impact of carbon dioxide flow rate on magnesium carbonate precipitation -step c(II) As outlined above, earlier experiments for step c(II) used a fixed CO2 flowrate of 20LPM, which established that -70-80% of magnesium in the calcium-depleted brine was removed under these conditions. Therefore, having optimised the pH setpoint for magnesium carbonate precipitation, it was next investigated how the flowrate of CO2 impacted the amount of solid produced and the consumption of NaOH. In accordance with the general methodology for step c(II) above, 25% NaOH was added to the calcium depleted brine at a rate to maintain a pH set point of 9.3, whilst CO2 was added at a rate of 5, 10, 15, 20 or 25 LPM (conditions 1-5, respectively). The measured change in concentration of magnesium in the overflow (post precipitation), and filtrate (from the precipitation tank solids), alongside the sodium hydroxide consumption, is reported below. Change in magnesium concentration and NaOH consumption with varying CO2 flow rates Condition Magnesium decrease in overflow (%) Magnesium decrease in fdtrate (%) NaOH consumption (kg / m3) 1 (5 LPM CO2) 10 (n=l) 0 (n=l) 1.2 (n=l) 2(10LPMCO2) 17 (n=l) 18 (n=l) 1.88 (n=l) 3 (15LPMCO2) 37 (mean, n=2) 36 (mean, n=2) 2.96 (mean, n=2) 4 (20 LPM CO2) 54 (mean, n=2) 59.5 (mean, n=2) 3.34 (mean, n=2) 5 (25 LPM CO2) 43 (n=l) 52 (n=l) 6.16 (n=l) As can be observed, CO2 flow rates lower than 20LPM indicated that there was a significant decrease in the amount of magnesium precipitated from the brine. Interestingly, at the CO2 flowrate of 25 LPM, the amount of magnesium precipitated decreased to less than 50%. NaOH consumption decreased at the lower CO2 flowrates, which was expected as less NaOH would be needed to react with the lower volume of CO2 that was added at the lower flowrates. At 25 LPM, the NaOH consumption became excessive (>6kg / m3) without the added benefit of an increased magnesium carbonate precipitation over 20 LPM conditions. Analysis of the precipitation tank solids and yield recovered from 1 L of brine is reported below in Figure 15. ICP-OES and XRD analysis of solid composition and yield of solid recoveredfrom 1L of brine (Figure 15) Condition Yield of solid ||||||||||||||| Mg % Nesquehonite (MgCO; 3H2O) % Monohydrocalcite (CaCO? H2O) % 1 (5 LPM CO2) 1.34 95 96.69 3.31 2 (10 LPM CO2) 1.98 95 95.75 4.25 3 (15 LPM CO2) 3.21 97 98.17 1.83 4 (20 LPM CO2) 6.28 98 98.63 1.37 5 (25 LPM CO2) 7.16 97 97.16 2.84 XRD and ICP-OES revealed that all CO2 flow rates investigated, magnesium carbonate was obtained in excellent purity (greater than 95%), with the highest purity of 98% at 20 LPM CO2. Furthermore, the yield of magnesium carbonate increased with increased CO2 flow rate, with over 6g obtained at flow rates greater than 20LPM. XRD analysis of the solids also proved that the magnesium carbonate was predominantly Nesquehonite (NQ). Overall, this experiment suggested that the most preferred flow rate for CO2 addition in step c(II) is 20 LPM. This is equivalent to the addition of CO2 at a flow rate of 2 L CO2 / L Brine. Investigating precipitation of magnesium hydroxide from calcium carbonate and magnesium carbonate depleted brine - step c(I) following step d(II) In a final experiment, the precipitation of magnesium hydroxide was investigated using calcium carbonate and magnesium carbonate depleted brine (i.e. performing steps c(I) and d(I) after steps a(I)-(III) and c(II)-d(II)). This was performed as a direct comparison to precipitation of magnesium hydroxide using calcium carbonate depleted brine (i.e. performing steps c(I) and d(I) directly after steps a(I)-(III)), with the difference in the brine compositions summarised below. Composition of CaCOs-depleted brine compared to CaCOs / MgCOs-depleted brine. Property ('at 'Os-depleted brine CaClh MgCth- depleted brine pH 8.95 9.44 Electrical Conductivity at 25 A°C (ApS cm) 78150 79200 Total Dissolved Solids (Calc.) (mg / L) 50817 51450 Total Hardness as CaCCh (mg / L) 9563 6080 Hydroxide Alkalinity as CaCCh (mg / L) 0 0 Carbonate Alkalinity as CaCCh (mg / L) 905 2535 Bicarbonate Alkalinity as CaCCh (mg / L) 655 928 Total Alkalinity as CaCCh (mg / L) 1560 3465 Sulfate as SO4 - Turbidimetric (mg / L) 4338 4350 Chloride (mg / L) 31700 34300 Calcium (mg / L) 85 21 Magnesium (mg / L) 2273 1465 Sodium (mg / L) 18700 20550 Potassium (mg / L) 810 789 Fluoride (mg / L) 1 1 Bicarbonate concentration (mol / L) 0.0131 0.0186 Carbonate concentration (mol / L) 0.0091 0.0253 Hydroxide concentration (mol / L) 0 0 Notably, the CaCCh / MgCCh-depleted brine had a higher pH, as well as a higher total alkalinity / TIC content than the CaCCh-depleted brine. This was consistent with the increased alkalinity in the overflow brine post-MgCCh precipitation, which was required to ensure a sufficient yield of nesquehonite. The total inorganic carbon concentration for the CaCCh / MgCCh depleted brine was double that of the CaCCh-depleted brine. Furthermore, it should be noted that the concentration of the magnesium in the CaCCh / MgCCh-depleted brine was almost half that of the CaCCh -depleted brine due to the precipitation of nesquehonite. As a final point, the CaCCh / MgCCh-depleted brine had a higher carbonate concentration than the CaCCh-depleted brine, which was also consistent with the brine having a higher pH (9.44 vs 8.95). This experiment was conducted in accordance with general procedure c(I) using 25% NaOH, wherein the pH for magnesium hydroxide precipitation was 10.8 to 11. A summary of the change in magnesium concentration in the brine is summarised below in Figure 16. Concentration of magnesium in brine compared to overflow (Figure 16) Condition Mg in brine (mg / L) Mg in overflow (mg / L) % change 1: CaCOs-depl eted brine, day 1, pH=10.8 2330 157 93 2: CaCO3-depl eted brine, day 2, pH=ll 2330 93 96 3: CaCO3 / MgCO3-depleted brine, day 1, pH=10.8 1470 61 96 4: CaCO3 / MgCO3-depleted brine, day 1, pH=ll 1460 28 98 Specifically, it can be seen in Figure 16 that all Mg(0H)2 precipitation days had a change in magnesium concentration between the brine and overflow of greater than 93%, which indicated excellent precipitation yields. Notably, Day 2 of Mg(0H)2 precipitation at pH 11 using CaCCh / MgCCh-depleted brine had the least residual Mg in the overflow (28 mg / L), and <5 mg / L in the filtrate from the precipitation tank solids. The consumption of NaOH for each day of the respective brines was very similar, when comparing it as kg of NaOH per kg of Mg(0H)2 precipitated. The molar composition of each solid isolated from the precipitation tank of conditions 1-4 above is reported below. Precipitation tank solids molar composition (%) Condition Fluoride Potassium Sodium Magnesium Calcium Chloride Sulfate 1 0.1 0.0 0.5 96.5 2.3 0.4 0.2 2 0.1 0.0 0.3 97.0 2.2 0.3 0.2 3 0.1 0.0 0.0 99.2 0.7 0.7 0.1 4 0.1 0.0 0.0 99.2 0.6 0.6 0.1 The precipitation tank solids for all samples indicated high magnesium content (greater than 96%). It should be noted that the samples precipitated from CaCO3-depleted brine had a calcium content of 2-3% merely because the brine started with a higher calcium content. The yield of solid filtered from 1 L of brine is reported below in Figure 17. Theoretical yield (based on starting Mg content in brine) compared to isolated yield (g solid fdteredfrom IL brine (Figure 17) Condition Theoretical yield (g / L) Isolated yield (g / L) 1: CaCOs-depleted brine, day 1, pH=10.8 5.96 5.43 2: CaCOs-depleted brine, day 2, pH=ll 7.35 5.60 3: CaCO3 / MgCO3-depleted brine, day 1, pH=10.8 2.24 3.42 4: CaCO3 / MgCO3-depleted brine, day 1, pH=ll 3.508 3.476 Notably, the yield for this step was very high, exceeding 3g / L even when using the CaCCL / MgCCL depleted brine. The potential benefit for this is that in a plant processing 3000 m3 / day, an estimated 9-10 tonnes (dry) of very pure Mg(0H)2 could be produced per day. ICP-OES analysis of the Mg(0H)2 filter press solids is reported below in Figure 18a and 18b. ICP-OES analysis of fdter press solids (Figure 18a and 18b) Condition CaCOi (%) Mg(OH)2 (%) NaCl (%) NaSO4(%) KC1 (%) Moisture content (%) 1 3 77 18 1 1 77.7 2 3 85 10 1 0 72.6 3 3 82 12 1 0 70.1 4 1 85 12 1 0 69.9 5* 4 95 1 0 0 79 *filter press solid from condition 1 repulped by mixing in tap water The ICP-OES analysis revealed that the NaCl content was high for all filter press solids under conditions 1 to 4 (>10%) and all precipitated solids had high moisture (>70%). The solids from condition 1 had flocculant added to aid settling, however, this proved to inhibit settling and increased moisture / NaCl content of the solid at 77.7% and 18 % respectively. However, as found in earlier experiments, the purity of Mg(0H)2 can be improved through repulping by mixing in tap water (condition 5, Figure 18b). By this method, the filter press solid from condition 1 above was repulped by mixing in tap water, which improved the Mg(0H)2 composition purity in the solid from 77% to 95%. This also greatly reduced halite contamination. XRD analysis was consistent with the ICP-OES data above for conditions 1-5. The particle sizes for all Mg(0H)2 precipitates was uniform, ranging from 9.5 to 12.7 pm (D[4,3]) which was similar to the results in Example 2, figure 9. Taken together, this experiment demonstrated the viability of sequential precipitation of calcium carbonate, followed by magnesium carbonate and then magnesium hydroxide. Furthermore, this experiment confirmed pH 10.8 to 11.0 is preferred for precipitation of magnesium hydroxide. Other variations to the disclosed embodiments can be understood and effected by those skilled in the art in practicing the claimed invention, from a study of the drawings, the disclosure, and the appended claims. In the claims, the word "comprising" does not exclude other elements or steps, and the indefinite article "a" or "an" does not exclude a plurality. The mere fact that certain measures are recited in mutually different dependent claims does not indicate that a combination of these measures cannot be used to advantage.

Claims

1. Process for producing calcium carbonate and either magnesium hydroxide or magnesium carbonate, using an aqueous brine solution which contains calcium and magnesium, comprising the steps of:a. (I) adding an alkalinity source to the aqueous brine solution such that the pH of the solution is between 8.7 and 9.3;a. (II) adding carbon dioxide gas to the aqueous brine solution produced in step a(I) such that the pH of the solution is between 5.0 and 6.0;a. (Ill) adding an alkalinity source to the aqueous brine solution produced in step a(II) such that the pH of the solution is between 8.5 and 9.5 and calcium carbonate precipitates;b. removing calcium carbonate produced in step a(III); then eitherc. (I) precipitating magnesium hydroxide by adding an alkalinity source to the calcium-depleted aqueous brine solution such that the pH of the solution is between 10.5 and 11.5 and magnesium hydroxide precipitates; andd. (I) removing magnesium hydroxide produced in step c(I); orc. (II) precipitating magnesium carbonate by adding an alkalinity source and carbon dioxide gas to the calcium-depleted aqueous brine solution such that the pH of the solution is between 9.0 and 10 and magnesium carbonate precipitates; andd. (II) removing magnesium carbonate produced in step c(II).

2. Process according to claim 1, wherein steps c(I) and d(I) are carried out.

3. Process according to claim 1, wherein steps c(II) and d(II) are carried out, optionallyfollowed by steps c(I) and d(I) using the calcium / magnesium-depleted aqueous brine solution.

4. Process according to claim 3, wherein step c(II) comprises simultaneously adding carbon dioxide gas and an alkalinity source to the calcium-depleted aqueous brine solution produced in step a such that the pH of the solution is between 9.0 and 10.0 and magnesium carbonate precipitates.

5. Process according to any preceding claim, wherein the aqueous brine solution is waste brine from a desalination plant.

6. Process according to any one of claims 1 to 4, wherein the aqueous brine solution is sea water.

7. Process according to any one of claims 1 to 4, wherein the aqueous brine solution isindustrial or municipal wastewater.

8. Process according to any one of claims 1 to 4, wherein the aqueous brine solution is a saline ground water.

9. Process according to any one of claims 1 to 4, wherein the aqueous brine solution is waste brine from a reverse osmosis plant or thermal desalination plant.

10. Process according to any preceding claim, wherein the alkalinity source is an alkali metal hydroxide, preferably sodium hydroxide.

11. Process according to any preceding claim, wherein the pH increase in step a(I) is to a value between 8.9 and 9.1, preferably to about 9.0.

12. Process according to any preceding claim, wherein the pH decrease in step a(II) is to a value between 5.2 and 5.6, preferably to about 5.4.

13. Process according to any preceding claim, wherein the pH increase in step a(III) is to a value between 8.8 and 9.5, preferably between 9.0 and 9.5, more preferably between 9.0 and 9.2, most preferably to about 9.1.

14. Process according to any preceding claim, wherein the pH increase in step c(I) is to a value between 10.8 and 11.2, more preferably between 10.8 and 11.0, most preferably to about 11.0.

15. Process according to any preceding claim, wherein the pH adjustment in step c(II) is to a final value between 9.3 and 9.7, preferably between 9.3 and 9.5.

16. Process according to any preceding claim, wherein adding carbon dioxide in step a(II) is at a flow rate of between 0.54 and 1.00 L CO2 / L Brine.

17. Process according to any preceding claim, wherein adding carbon dioxide in step c(II) is at a flow rate of 2 L CO2 / L Brine.

18. Process according to any preceding claim, wherein the alkalinity source of step a(I) is lime.

19. Process according to any preceding claim, wherein the removal of calcium carbonate in step b, and / or the removal of magnesium hydroxide in step d(I) or the removal of magnesium carbonate in step d(II) comprises the steps of:(i) allowing the calcium carbonate and / or magnesium hydroxide or magnesium carbonate precipitate to settle;(ii) removing the supernatant liquid; and(iii) filtering the remaining slurry.

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