CA2915623A1 - Separation of carbon dioxide from flue gases - Google Patents

Abstract

A method of separating carbon dioxide from a flue gas includes exposing the flue gas comprising carbon dioxide to salt water to generate a solution, monitoring pH
of the solution, when a pH condition is met, adjusting the pH to precipitate a solid carbonate material from the solution, separating the solid carbonate material from the solution, and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

Description

SEPARATION OF CARBON DIOXIDE FROM FLUE GASES
Technical Field [0001] The present invention relates to the separation of carbon dioxide from flue gases for sequestering the carbon dioxide.
Background

[0002] Sequestration of carbon dioxide as solid carbonates is desirable to reduce the levels in or released to the atmosphere. Mineral carbonation, which is the process of capturing the CO2 in the atmosphere in the form of solid carbonates through the reaction of CO2 with silicates, is a spontaneous, thermodynamically favourable process. Solid carbonates are a long-term geological sink of carbon and thus their formation is desirable. Unfortunately, the kinetics of natural mineral carbonation is very slow and the process is only significant over geological time periods, i.e., millions of years.

[0003] Accelerated formation of solid carbonates is observed in biological systems, particularly in corals, bivalve molluscs, echinoderms, and foraminifera.
These organisms have developed mechanisms to induce and accelerate the precipitation of carbonates, which is required for their skeletons, in natural saline waters. Mimicking such biological approaches to achieve the accelerated formation of solid carbonates is desirable.

[0004] Precipitation of carbonates in seawater does not occur spontaneously.
Economically feasible processes to induce and accelerate precipitation are desirable.
Summary

[0005] According to an aspect of an embodiment, a method of separating carbon dioxide from a flue gas is provided. The method includes exposing the flue gas comprising carbon dioxide to salt water to generate a solution, monitoring pH
of the solution, when a pH condition is met, adjusting the pH to precipitate a solid carbonate material from the solution, separating the solid carbonate material from the solution, and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

[0006] According to another aspect of an embodiment, a method of separating carbon dioxide from a flue gas includes exposing an underground silicate deposit to hydrochloric acid to generate a brine, recovering the brine at surface, exposing the flue gas comprising carbon dioxide to the brine, monitoring pH of the brine, when a pH condition is met, adjusting the pH, precipitating a solid carbonate material from the brine, separating the solid carbonate material from the brine, and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

[0007] According to yet another aspect of an embodiment, a method of separating carbon dioxide from a flue gas includes exposing an underground silicate deposit to hydrochloric acid to generate a brine, recovering the brine at surface, exposing the flue gas comprising carbon dioxide to the brine, introducing ammonia to the brine to generate an ammonium salt and a solid carbonate material, precipitating the solid carbonate material from the brine, separating the solid carbonate material from the brine, and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

[0008] According to still another aspect of an embodiment, a method of separating carbon dioxide from a flue gas includes exposing the flue gas comprising carbon dioxide to salt water to provide a solution, monitoring pH of the solution, adjusting the pH to generate a basic solution and precipitate a solid carbonate material from the solution, separating the solid carbonate material from the solution, and storing the solid carbonate material to inhibit carbon dioxide from entering the atmosphere.

[0009] According to a further aspect of an embodiment, a method of sequestering carbon dioxide includes exposing a flue gas comprising carbon dioxide to sea water to provide a solution, monitoring the pH of the solution, when a pH
condition is met, adjusting the pH to precipitate a solid carbonate material from the solution, and storing the solid carbonate material in the sea water to inhibit carbon dioxide from entering the atmosphere.

Brief Description of the Drawings

[0010] Embodiments of the present invention will be described, by way of example, with reference to the drawings and to the following description, in which:

[0011] FIG. 1 is a schematic diagram of a system for the capture of CO2 using sea water and subsequent neutralization of the acidified sea water according to an embodiment;

[0012] FIG. 2A is chart showing an equilibrium and operating diagram for absorption with sea water at variable pH;

[0013] FIG. 28 is chart showing an equilibrium and operating diagram for absorption with sea water at constant pH = 8.2;

[0014] FIG. 3 schematic diagram of an electrochemical cell for the regeneration of sea water acidified after contact with flue gas according to an embodiment;

[0015] FIG. 4 is a chart showing the energy utilized for electrolytic mineralization of CO2 according to an embodiment;

[0016] FIG. 5 is a schematic diagram of a system for CO2 mineralization with silicates according to another embodiment; and

[0017] FIG. 6 is a schematic diagram of a system for CO2 mineralization with ammonia and ammonia regeneration according to another embodiment.
Detailed Description

[0018] For simplicity and clarity of illustration, reference numerals may be repeated among the figures to indicate corresponding or analogous elements.
Numerous details are set forth to provide an understanding of the examples described herein. The examples may be practiced without these details. In other instances, well-known methods, procedures, and components are not described in detail to avoid obscuring the examples described. The description is not to be considered as limited to the scope of the examples described herein.

[0019] The disclosure generally relates to a system and a process for separating carbon dioxide from a flue gas. The includes exposing a flue gas that includes carbon dioxide to salt water to generate a solution, monitoring pH of the solution, adjusting the pH to precipitate a solid carbonate material from the solution when a pH condition is met, separating the solid carbonate material from the solution; and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

[0020] Several solvents have been previously proposed and tested for CO2 capture. The use of some solvents such as amines or ionic liquids, however, increases the overall cost of carbon sequestration because of the cost of the solvents and the large amount of energy utilized for regeneration of the solvents.
Water may be utilized for capturing carbon dioxide from flue gas, however the equilibrium concentration of CO2 in water is relatively low. Nonetheless, there is a growing interest in using water, and in particular, particularly sea water, as a medium for capturing and sequestering CO2.

[0021] The supply of cations utilized for carbonate precipitations that exist within sea water makes sea water a desirable. The dissolution of carbon dioxide in water is generally believed to follow the mechanism shown below:
CO2 (g) Ã4 CO2 (aq) AH = ¨19.75 kJ/mol (1) r CO2 (am H2O 4-+ H2CO3 AHr = ¨0.56 kJ/mol (2) CO2+ OH- . HCO3- ATP= ¨46.655 kJ/mol (3) r H2CO3 4--* HCO3- + H Aii -= 9.72 kJ/m o I (4) r HCO3- 4-4 C032- H Air =14.7 kJ/mol (5) r where the enthalpies of reaction at standard conditions were determined using data from Dean, IA. Lange's handbook of chemistry, 15th Edition, 1999.

[0022] The overall reaction is:
CO2 (g) H2O 4-- C032- 2H AHr = 4.11 kJ/mol, Aq=103.6kJ/mol (6) or at higher pH, CO2 (g) OH-. C032- + H Aiir = ¨51.7 kJ/mol, Aq= 23.77 kJ/mol (7).

[0023] The standard Gibbs energy of reaction for the global reactions (6) and (7) is positive, indicating that the conversion of CO2 into carbonate is not spontaneous at 25 C. Reaction (7), however, may be spontaneous at lower temperatures because this reaction is exothermic.

[0024] If divalent ions are present in solution, e.g., Ca2+, Mg2+, Sr2+, the formation of solid carbonates has the potential to occur. For example, calcium ions react with carbonate ions in solution to yield calcite or aragonite:
Ca2+ + C032- CaCO3 (calcite or aragonite) (8).

[0025] By combining reactions (6) or (7) with reaction (8), two global reactions are possible:
CO2 (g) + H20 + Ca2+ 4-> CaCO3 (calcite) + 2 H+ A11r =14.7kJ/mol, Aq= 56.0 kJ/mol (9) CO2 (g) OW + Ca2+
CaCO3 (calcite) H+ AHr = -41.1k3/mol, AGr = -24.7 kJ/mol (10)

[0026] From the values of the enthalpy of reaction and Gibbs energy of reaction, it is clear that reaction (9) is not spontaneous at any temperature, while reaction (10) is spontaneous at all temperatures. This means that spontaneous precipitation of calcite is only spontaneous at high pH. The same pattern holds for the other carbonates (i.e. aragonite, magnesite, and dolomite).

[0027] The formation of carbonates may theoretically take place when CO2 is solubilised in a solution containing divalent cations resulting in the precipitation of several different solid carbonates including calcite, aragonite, dolomite, and magnesite. In practice, however, formation of solid carbonates is rarely observed.
In fact, sea water is often supersaturated with carbonates. As an example, Dolomite formation has not been observed in sea water, even though the ionic product of calcium, magnesium, and carbonate is several orders of magnitude greater than the solubility product of dolomite. Sedimentary deposits, however, are rich in magnesium-calcite and dolomite, indicating that at some time during the earth's geological history precipitation of carbonates has occurred on a very large scale as described in, for example, Burns, S.J. et. al. (2000) Dolomite formation and biogeochemical cycles in the Phanerozoic. Sedimentology 47(1), 49-61 (hereinafter "Burns").

[0028] The prior art related to carbonate precipitation in sea water suggests that solid carbonate formation is mediated by the presence of living organisms that either change the redox potential of the water, as described in, for example, Burns, or by the raise the local concentration of ions through the use of biological membranes, as described in, for example, Al-Horani, F.A. et al (2003) The mechanism of calcification and its relation to photosynthesis and respiration in the scleractinian coral Galaxea fascicularis. Marine Biology 142, 419 - 426 (hereinafter "Al-Horani").

[0029] The precipitation of calcium and magnesium carbonates in sea water (pH = 8.2) is catalyzed by several living forms that use those carbonates to build their skeletons (i.e. corals, bivalves mollusca, echinoderms, and foraminifera).
These organisms utilize different strategies to carry out and accelerate carbonate precipitation. In the case of corals, calcium transporters are responsible for pumping calcium ions against their concentration potential causing an increase in the oversaturation of calcium (though sea water is already supersaturated in calcium). The energy to drive the calcium transporters is provided by photosynthesis, as described in, for example, Al-Horani.

[0030] When a gas with a high concentration of CO2, such as for example flue gas, is put in contact with sea water, CO2 will solubilize in the water increasing the concentration of soluble carbon species in the sea water, as well as the concentration of protons (H+). The increased acidity, however, impedes the precipitation of solid carbonates as shown in the above Equations (9) and (10).
Typically, the solubility of carbonates increases as pH decreases. Once the acidified water re-enters the environment and reaches equilibrium with the predominant concentration in the atmosphere, most of the solubilized CO2 will have come out of solution and been released into the atmosphere, resulting in a zero net carbon sequestration. Thus, challenges exist to carbon sequestration through carbon solubilizing in sea water. To achieve a positive net carbon sequestration using sea water, solid carbonates must be formed and removed prior to returning the water to the sea.

[0031] To illustrate the effect that CO2 dissolution has on the solubility of solid carbonates in sea water, and how to enhance carbonate precipitation, the thermodynamic equilibrium of the CO2-carbonate-sea water system was simulated using the geochemical software PHREEQC version 3Ø PHREEQC is a computer program for speciation, batch-reaction, one-dimensional transport, and inverse geochemical calculation available from the U.S. Geological Survey and is described more fully at, for example, Parkhurst, D.L., and Appelo, C.A.J., 2013.
Description of input and examples for PHREEQC version 3 ¨ A computer program for speciation, batch-reaction, one-dimensional transport, and inverse geochemical calculations:
U.S. Geological Survey Techniques and Methods, book 6, chap. A43, 497.

[0032] A gas stream representing flue gas is assumed to contact a liquid stream representing sea water saturated with CO2 and oxygen at the predominant atmospheric concentration of these gases. Flue gas is the product of combusting a hydrocarbon fuel with air. Normally, air is supplied in excess of the stoichiometric requirement in order to ensure complete combustion. The molar composition of the flue gas resulting from burning natural gas (CH4), assuming a 15% excess air, is shown below:
Component N2 02 CO2 Ar H20 % mol 71.67 2.88 8.21 0.86 16.35

[0033] The composition of the sea water stream used in the simulation is reported in Table 1 of Appendix 1 to this disclosure. The results from the PHREEQC
simulation are presented in Appendix 2 to this disclosure. The final pH of the sea water solution is about 5.8 after being exposed to flue gas and reaching equilibrium conditions, as shown in part B of Appendix 2. This acidified solution is then neutralized with a 0.5M sodium hydroxide solution to bring the pH close to its original value (-8.2).

[0034] As set out in Appendix 2, the saturation indexes for different carbonates are presented for the original sea water solution, for the acidified sea water solution, and after neutralization. The saturation index (SI) is defined as:
SI = log (¨IAP (11) Ksp where, TAP is the ion activity product and lc is the equilibrium solubility constant.

[0035] Positive SI values indicate that the IAP is greater than the K51, and precipitation should occur. For negative SI values the IAP is smaller than the K51, and precipitation is inhibited. Values near zero indicate equilibrium. For calcite, SI
values around 0.2 log units may suggest equilibrium if pH measurements are considered to be reliable. The activity of the carbonate component (C032-) is a direct function of pH: a decrease in pH of one unit will decrease the activity of C032 byone log unit, as described at, for example, Ouled Ameur, Z. et. al. (2015) Stimulation of High Temperature SAGD Producer Wells Using a Novel Chelating Agent (GLDA) and Subsequent Geochemical Modeling Using PHREEQC. SPE
International Symposium Oilfield Chemistry, The Woodlands, Texas, USA, 12-15 April, 2015.

[0036] As shown in the PHREEQC output file in Appendix 2, sea water is initially oversaturated with calcite, aragonite, and dolomite. This oversaturation is a characteristic of sea water that has been extensively documented, such as in, for example, Mitchel, M.J. et. al. (2010) A model of carbon dioxide dissolution and mineral carbonation kinetics. Proceedings of the Royal Society A 466, 1265-1290.

[0037] After equilibrium with flue gas, the sea water is no longer oversaturated with carbonates. The increased partial pressure of CO2 in the flue gas causes a drop in the pH of the sea water, increasing the solubility of carbonates. After the pH of the acidified sea water has been brought back to the original sea water pH by neutralization with NaOH, the saturation indexes for calcite (1.26), aragonite (1.11) and domolite (3.40) are greater than one, indicating that precipitation of dolomite, calcite, and aragonite is, at least theoretically, possible.

[0038] As described above, a prerequisite for precipitating carbonates out of sea water is to counter the excess of hydronium cations (W). Simple neutralization with, for example, a base (hydroxide ion donor) is effective at shifting the equilibrium to the point at which precipitation of carbonate is feasible.
Alternatively or additionally, a buffer solution may also be used to maintain a constant pH
in the solution. Alternatively, or additionally, neutralization may be performed by removing H ions using a membrane. Continued movement of the H ions across the membrane may be facilitated through a reaction that consumes protons on one side of the membrane.

[0039] Fig. 1 shows a schematic diagram of an example system 100 for the capture of CO2 using sea water and subsequent neutralization of the acidified sea water. The system includes a CO2 absorption column 102, into which sea water 104 and flue gas 106 enter. Within the CO2 absorption column 102, CO2 from the flue gas 106 is absorbed by the sea water 104. Treated gas 108 and acidified sea water 110 exit the CO2 absorption column 102. The acidified sea water 110 enters a neutralizer 112, which removes H+ ions 114 from the acidified sea water 110 to water's increase the pH. Neutralized sea water 116 enters a separator 118 in which CaCO3 precipitate 120 is removed, leaving spent sea water 122.

[0040] The CO2 absorption column 102 may be designed taking into consideration the equilibrium between CO2 in air and CO2 in sea water, as well as the kinetics of CO2 dissolution and hydration. It is possible to carry out the absorption at constant or variable pH. To maintain a constant pH, it is necessary to remove the excess hydronium ions formed by the dissolution of CO2. If the produced hydronium are not removed or otherwise neutralized, then the pH will be lower as the sea water moves along the absorption column 102. The equilibrium and proposed operating diagram 200 corresponding to a variable pH operation for the absorption column is shown in Figure 2A.

[0041] Operation at variable pH requires a very high liquid to gas ratio.
As shown in FIG. 2A, the minimum molar L'/G' ratio is 1770 mol H20/ mol air, which is approximately equivalent to 1050 kg sea water /kg of flue gas. The calculated minimum liquid to gas ratio at variable pH is much higher than the economical optimum calculated in previous section. As it is not feasible to operate the column at a liquid to gas ratio lower than the minimum (because it will require the operating line to cross the equilibrium line), then it is concluded that the absorption will have to be conducted at a ratio much higher than the economical optimum.
Alternatively, the absorption may be carried out at constant pH, which effectively modifies the equilibrium curve 200, as illustrated in chart 202 shown in the Figure 2B.

[0042] At pH 8.2, the minimum liquid to gas ratio is only 7.8 mol H2O /
mol of air, or 4.6 kg of sea water per kg of flue-gas. Operating at 1.5 times the minimum L/G ratio, or about L'/G' = 12.4, eight ideal separation stages are needed to achieve the proposed absorption. If the operation is performed at the economical L/G ratio, near 60 kg of sea water per kg of flue-gas or L'/G' = 100, the number of ideal equilibrium stages required may be reduced to two. If only a reduced number of separation stages are required, then it might be feasible to perform inter-stage neutralization of the excess acid.

[0043] The size (height and diameter) of the absorption column 102 absorbing the CO2 into sea water may be determined by considering the mass transfer and reaction kinetics limitations. The CO2 hydration reaction is usually considered to be the limiting reaction step in the CO2 dissolution and mineralization process, while the other reactions are assumed to occur almost instantly. It must be noted, however, that the carbonate precipitation and crystallization reactions are indeed mass-transfer limited. For the purpose of this preliminary estimation, it will be assumed than only the hydration reaction limits the whole process, and all other species reach equilibrium instantly.

[0044] From reaction (2), the rate of disappearance of CO2 may be written in terms of the forward and reverse reaction constants and the concentration of the species intervening in the reaction. Assuming that instantly the equilibrium between the CO2 in the gas phase and the CO2 in the liquid phase is reached, as well as the equilibrium between carbonic acid and the bicarbonate and hydrogen ions, then it is possible to write:
r rc02 = [CO 2(aq)]+ k2- [14 2CO3 = ¨k;Keqi [CO2( ]-1- 2 PC031141 (Si) Keq3

[0045] The forward reaction is assumed to have a rate constant k2+ = 0.06 (see, for example, Bond, G.M. et al (2001) CO2 Capture from coal-fired utility generation plant exhausts, and sequestration by a biomimet route based on enzymatic catalysis - Current Status. NETL, First National Conference on Carbon Sequestration), while the backward reaction is assumed to have a rate constant k2-= 20 s-1 (see, for example, Mitchel, M.J. et al (2010) A model of carbon dioxide dissolution and mineral carbonation kinetics. Proceedings of the Royal Society A
466, 1265-1290).

[0046] If each single stage is treated as a Continuous Stirred Reactor (CSTR) (ideal mixing in the stage), then the stage volume may be calculated as:

V CSTR
(52) r rco, 'exit

[0047] The maximum possible conversion is determined by the equilibrium conditions. At the equilibrium (with X (reaction advancement at anytime) = Xeq (reaction advancement at equilibrium)), the reaction rate is effectively zero, which will require an infinite reactor volume to occur. To calculate the required volume for a conversion less than Xeq, the concentration of all relevant species leaving the reactor may be determined, as the conditions inside the reactor (or equilibrium stage) are assumed to be the same as those of the stream leaving the reactor (i.e.
ideal mixing). The concentrations of all species at equilibrium may be first calculated and then the concentrations at the specified conversion point may be determined by interpolation. The reaction rate was then calculated using Eq.
(Si) and the concentrations at the specified conversion point. The equilibrium conversion between sea water and a typical power plant flue gas (usually containing 8% of CO2) at atmospheric pressure was determined to be roughly 60% (see, for example, Xu, Xiaochun et al. (2003) Separation of CO2 from Power Plant Flue Gas Using a Novel CO2 "Molecular Basket" Adsorbent, Fuel Chemistry Division Preprints 2003, 48(1), 162).

[0048] Table 2 included in Appendix 3 to this disclosure shows the estimated reactor volume for different conversions as well as the expected reactor volume if a catalyst is used to speed up the reaction. In one case, it is assumed that the catalyst may provide a 70-fold increase in reaction rate, similar to what is expected if the carbonic anhydrase enzyme is used. In the second case, a 14-fold increase in the reaction rate is considered, which corresponds to what is expected if nickel nanoparticles are used (see, for example, Ouled Ameur, Z. and Husein, M.M.

Electrochemical behavior of Potassium Ferricyanide in aqueous and (w/o) Microemulsion in the Presence of Dispersed Nickel Nanopartciles. Separation Science and Technology 48(5): 681-689).

[0049] To determine the amount of H+ ions to be removed or neutralized from the acidified sea water to cause precipitation of CaCO3, a total mass balance and elemental balances, satisfying all equilibrium conditions, may be performed around the process envelope, such as the envelop shown as the dashed line 124 in the example system 100 shown in FIG. 1.

[0050] The total mass balance and elemental balances are given by:
Total mass balance is given by:
Go +Lc,(12a) = + mcaco, Carbon balance is given by:
G0[CO2]0 +L1(CO*2]0 +[HCOs]o +[C032]0)= G1[CO2]1 +Li([CO*2], +[HCO]i +[0:X1)+ncaco, (12b) Hydrogen balance is given by:
L00-1C0i]0 +[11-]0+ 2 [H20]3 + [OH-]= (HCOi + [H+ + 2 [H20]1 + [OH-]1 )+ nw_ (12c) Oxygen balance is given by:
2G0[CO2]0+ Lo picoijo + [H20j0 + [OH- ]0 2 [CO*2 + 3[C03210 =
2G0 [CO2 Jo + (3 [HCOi +[H20]1 + [OH-]1 + 2 [CO*2 + 3 [C032 -]
)+ 3ncaco, (12d) where:
Lo is the initial sea water mass flow rate in kg/[t]
Go is the initial flue gas conditions mass flow rate in kg/[t]
L1 is the acidified sea water mass flow rate (after reaching equilibrium) in kg/[t]
G1 is the flue gas flow rate (at final conditions) in kg/[t]
IT is the mass flow rate of protons ions in kg/[t]
mCaCO3 is the mass flow rate of calcium carbonate in kg/[t]
n CaCO3 is the molar flow rate of calcium carbonate in mole/[t]
and nu+ is the molar flow rate of proton ions in mole/[t]

[0051] The equilibrium relations are given by:

[CO*2] = IC, fcc,2 (13a) [HCC30][H'] = K1[CO*2]
(13b) [C0][H' ]= K2[HCOi]
(13c) [OH-][H]= K õ, (13d) [co;-][ca 2 5 Kfpalcite (13e) [ 3il a c02-c 2 5_ p KsAragonite (13f) [ ,CO2-i[mg p 2 < K smagnesite (13g) [co[ 2 [ca2 11 mg2 < KsDpolomite (13h)

[0052] The values for K0 and Kw may be obtained from, for example Dickson, A.G. and Goyet, C. 1994, Handbook of methods for the analysis of the various parameters of the carbon dioxide system in sea water, U.S. Department of Energy ORNL/CDIAC-74; the values for K1, K2 may be obtained from Mojica, F.J. and Miller, F.J. 2002, The values of pKi + pK2 for the dissociation of carbonic acid in sea water, Geochimica et Cosmochimica Acta 66(14), 2529-2540; the solubility product constants (Ksp) for calcite and aragonite may be obtained from Al-Rawajfeh, A.E.
(2004) Modelling and simulation of CO2 release in multiple-effect distillers for sea water desalination, Martin Luther Universitat Halle Wittenberg, PhD thesis, 2004;
and the values for all other constants may be obtained may be obtained from Dean.
The fugacity of CO2 may be assumed to be equal to its partial pressure (i.e., ideality may be assumed).

[0053]
Assuming a sea water inlet flowrate of 1 kg/[t] and a flue gas flowrate of 0.001 kg/[t], the composition and flowrate of the outlet streams may be calculated as shown in Table 3 included in Appendix 4 to this disclosure. In the calculations reported in Table 3, the amount of H+ produced is adjusted to give the treated gas a CO2 concentration equal to that of atmospheric air (around 0.039%).
That is, from a greenhouse gas perspective the discharged treated gas may be considered to be equal to atmospheric air (though depleted in oxygen).

[0054] For capturing 1 MM tons of CO2 per year, the following rates are desired:
CO2 sequestered = 32.15 kg/s = 116 t/h H+ consumed = 1787 molts CaCO3 (s) produced = 115 kg/s = 412 t/h Fresh sea water = 897626 m3/h = 249 m3/s

[0055] In an example in which caustic soda (NaOH) is used to neutralize the excess H+, utilizing 2.22 ton of NaOH facilitates capturing one ton of CO2. In this example, for a bulk cost for NaOH of $415/ton of NaOH, the cost for CO2 capture will be $922 per ton. Therefore, cost considerations may make neutralization utilizing NaOH unfeasible.

[0056] Precipitation of carbonates may be accomplished if a weak base is added to the acidified sea water, or to the sea water before CO2 absorption.
Ammonia (NH3) may be used in a similar way as it is used in the Solvay process for the production of sodium bicarbonate.

[0057] When ammonia is bubbled in water the following reaction takes place:
NH3 (g) H20 4-> NH4 + OH- AI/;? = -31.5 kYmol, AGr =17.1 kJ/mol (14) or at lower pH:
NH3 (g) 1-1+ 4--> NH4+ AH = -87.3 kJ/mol, AG = -62.8 kJ/mol (15)

[0058] If reactions (14) and (15) are coupled to reactions (6) and (7), the following global reactions may take place:
2NH3 + Ca2+ + CO2 + H2O-CaCO3(s) + 2 NH4 + AHr = -159.9 k3/mol, AG ¨69.6 kJ/mol (16) NH3+ Ca2+ + CO2 + H20 CaCO3(s) + H+ + NH4 + AHr = -72.6 kJ/mol, AG = -6.8 kJ/mol (17) NH3 + Ca2+ + CO2 + OH- CaCO3(s) + NH4 AH(r)= -128.4 kJ/mol, AG = -86.7 kJ/mol (18)

[0059] All three reactions (16 to 18) are exothermic and spontaneous. The pH
of the system will determine which reaction is dominant.

[0060] To determine the amount of ammonia that is desirable to achieve the mineralization of CO2, the hydrogen balance formulated in Eq. (12c) may be modified as follows:
I0(HC0i10 + [H 10 + 2 [H20]0 +[OH-10)+ =
41-1C0i]1 +[H ]1 +2[H20]1 2 [H20], + [OH-], +3[N113]1 + 4[N114],) (12c*) and the following equilibrium relation may be utilized to the set of equilibrium constraints:
[NH, ][OH ] = Kb [NH3]
(13i)

[0061] Solving the modified system of equations (12* and 13i), the amount of ammonia that facilitates the capture 1 ton of CO2 is calculated as 1.2 ton of NH3. At a current cost of $600/t of NH3 the cost of carbon capture (due to NH3 only) may be $718/ton of CO2. Therefore, cost considerations may make neutralization utilizing NH3 economically unfeasible. However, a closed-loop operation with ammonia regeneration may reduce operating costs and will be described later in this disclosure.

[0062] As an alternative to neutralization with a base, the excess hydronium ions (H+) may be removed from the acidified sea water using a membrane; for example a polymeric proton exchange like Nafion could be used for this purpose. An anionic membrane may also be used to allow negatively charged ions (mainly chlorides and fluorides) to move out of the acidified water in order to maintain electroneutrality. To sustain a continuous flow of protons across the membrane an electrolytic cell may be utilized in order to convert the protons into hydrogen and the chloride ions into chlorine.

[0063] An example of an electrochemical cell 300 that may be utilized to sustain a continuous flow of protons is shown in Fig. 3. The electrochemical cell 300 includes a water inlet 302 through which acidified sea water flows into the cell 300, and a water outlet 304 through which water exits the cell 300. The cell 300 includes a cathode 306 and an anode that are coupled via power supply 310. A
proton exchange membrane 312 separates the inlet 302 from the cathode 306, and an anion exchange membrane 314 separates the inlet from the anode 308. The cell 300 also includes a hydrogen outlet 316 through which hydrogen exits the cell 300, and a chlorine outlet 318 through which chlorine exits the cell 300.

[0064] The electrochemical reactions of halogens are highly reversible, as described in, for example, Srinivasan 5, and Kirby B. Status of fuel cell technologies. In: Fuel Cells: From Fundamentals to Applications, Editor: S.
Srinivasan. Springer, 2006. For this reason, there has been a great interest in developing regenerative fuel cells based on the H2/C12 system for energy storage.
Cell designs for energy storage applications desirably have a high energy conversion efficiency; therefore they are a desirable start point for developing an electrochemical cell for enhancing carbonate precipitation in sea water.

[0065] Although several designs have been proposed for the H2/C12 system, as described in, for example, Thomassen M, Sandnes E, Borresen B, and Tunold R.
(2006) Evaluation of concepts for hydrogen - chlorine fuel cells. Journal of Applied Electrochemistry 36, 813-819 (hereinafter "Thomassen"), the H2/Cl2 galvanic cell is still considered an immature technology, as discussed more fully in House K.Z., House C. H., Schrag D. P., and Aziz M. J. (2009) Electrochemical acceleration of chemical weathering for carbon capture and sequestration. Energy Procedia 1, 4953-4960.

[0066] Thomassen discloses the performance of an ordinary polymer electrolyte membrane fuel cell based on a Nafion membrane, a fuel cell based on a combination of circulating hydrochloric acid and a Nafion membrane, and a system based on a phosphoric acid doped Polybenzimidazole (PBI) membrane. Thomassen concludes that even though the three systems studied achieved open circuit voltages close to the reversible and exhibited fast electrode kinetics, stable operation was not possible due to electrocatalyst corrosion.

[0067] More recently, the use of a ruthenium alloy oxide [(Ru0.09000.90304]
electrocatalyst for the chlorine electrode has been considered, as described more fully in Huskinson B., Rugolo J., Mondal S.K., and Aziz M.J. (2012) A high power density, high efficiency hydrogen-chlorine regenerative fuel cell with a low precious metal content catalyst. Energy Environ. Sci. 5, 8690 - 8698.

[0068] In this example system Nafion 112 was used as the proton-exchange membrane. The ruthenium loading in the chlorine-side electrode was 0.15 mg Ru/cm2, while the hydrogen-side electrode used a standard ELAT gas diffusion electrode (GDE) with a Pt loading of 0.5 mg/cm2. The reported power density was approximately 0.4 W/cm2 at 90% galvanic efficiency.

[0069] To evaluate the feasibility of using an electrochemical cell for removing excess hydronium ions and to promote carbonate precipitation, it may be desirable to estimate the energy consumption associated with operating the cell. The electric potential, E, that facilitates operating an electrochemical cell may be determined using the Nernst equation:
E = E0 - RT In Q
zF
(19) where E is the standard cell potential, Q is the reaction quotient, z is the number of moles of electrons transferred in the cell reaction, R is the universal gas constant, F is the Faraday constant, and T is the absolute temperature.

[0070] The reaction quotient Q in Eq. (19) may be calculated assuming that the activity of the gas species involved in the electrochemical reaction is equal to their partial pressure (P,), and the activity of the ionic species is equal to their molar concentration in the solution.

[0071] For the example electrochemical cell 300 illustrated in Fig. 3, the global reaction may then be written as:
2 H+ + 2 C1 4-4 C12 H2 = -1.36 V
(20) and cell potential that facilitates the drive reaction (20) forward is estimated at E=-1.365 V. In practice, however, an overvoltage is desired to result in increased power consumption.

[0072] From the reduction half-reaction utilized in the example electrochemical cell 300 shown in Fig. 3, it may be seen that for each mole of H+
reduced, 1 mole of electrons is consumed. Therefore the electron consumption rate for sequestering 1MM ton of CO2/year will be 1787 mol

[0073] The power utilization may be given by:
P = lie FE = 1787 mol/s x 96484.56 C/mol x1.365 V 106 = 235 MW

[0074] 235 MW relates to the power utilization if no overvoltage is applied, to carry out the specified reaction. For an overvoltage of 0.8 V, the power consumption may be calculated as 373 MW.

[0075] The hydrogen gas produced may be used to generate electricity and partially off-set the power utilization of the process. Hydrogen may be converted to electricity through a combined cycle or in a fuel cell. Alkaline fuel cells have demonstrated efficiencies up to 60% (see, for example, Mulder, G. et al.
(2008) Market-ready stationary 6 kW generator with alkaline fuel cells. ECS
Transactions 12, 743-758), while phosphoric acid fuel cells, proton exchange membrane fuel cells, and solid oxide fuel cells may be up to 85% efficient when used in co-generation of electricity and heat (see, for example, Hamelin J., Agbossou K., Laperriere A., Laurencelle F, Bose T.K. (2001) Dynamic behavior of a PEM fuel cell stack for stationary applications. International Journal of Hydrogen Energy 26(6), 625-629, and Edwards P., Kuznetsov V., David W.I.F., Brandon N.P. (2008) Hydrogen and fuel cells: Towards a sustainable energy future. Energy Policy 36(12), 4356-4362). The efficiency of combined cycle gas turbines have increased in recent years and typical values of newly-build power plants are reported in around 61% (see, for example, Robb D. (2010) CCGT: Breaking the 60 per cent efficiency barrier. Power Engineering International 18(3)). Assuming a conversion efficiency of 70% the process will potentially produce 179 MW of power.
Consequently, the net energy utilized will be:
0 V overvoltage: Pnet = 56 MW = 1.7 GJ/ton CO2 = $29/ton CO2 0.8 V overvoltage: Pnet = 194 MW = 6.0 GJ/ton CO2 = $101/ton CO2, where electricity is assumed to cost 6 ct/kWh.

[0076] An alternative cell design aimed at reducing the net energy consumption, and in turn reducing the cost/ton of CO2 capture, consists of bubbling oxygen at the cathode, which suppresses hydrogen gas production. In this case, the half-cell reaction is:
02 (g) 4 H+ + 4e- *-4 2 H20 E = 1.229 V (21)

[0077] And the overall electrochemical reaction is:
4 H+ + 02(9) 4 a- 2 Cl2 + 2 H20 E = -0.131 V (22)

[0078] The energy utilized by this alternative cell design is:

E = E __ in 4F [1-Ff[C1-]2/33 - = -0.136 V (23)

[0079] Consequently, the net energy utilized will be:
0 V overvoltage: Pnet = 23.4 MW = 0.73 G3/ton CO2 = $12/ton CO2 0.8 V overvoltage: Pnet = 161 MW = 5.0 GJ/ton CO2 = $84/ton CO2.

[0080] Depending on the operating temperature of electrolitical cell, heat from the produced water or steam could be recovered or integrated to adjoining processes.

[0081] The calculations in the previous description were performed for an assumed sea water rate of 1 kg and 0.001 kg of flue gas, for a Lo/Go ratio equal to 1000. Modifying this ratio affects both the relative amount of sea water utilized and the energy utilized for pumping, as well as the final equilibrium points and the net energy utilized for electrolysis.

[0082] The electrochemical system described in this disclosure, such as the example electrochemical cell 300, may be similar in practice to existing membrane-based electrochemical cells for the chlor-alkali process. However, certain aspects of the carbonation process are different from the chlor-alkali processes, which may utilize a new cell design or special membranes. Among the aspects that a cell design may take into consideration are:
= Potential formation of precipitates inside the cell = Ion concentrations lower than in a conventional chlor-alkali system if untreated sea water is to be used = Presence of multiple contaminants that may adversely affect membrane performance

[0083] Cells may be designed to minimize or inhibit the potential formation of a solid layer on the membrane surfaces. Equipment may also be designed to allow recycling of acid effluent into the membrane chambers in case solids are deposited on the membrane.

[0084] In Table 4 included in Appendix 5 to this disclosure, the power utilization and sea water consumption per ton of CO2 captured are shown for different liquid to gas ratios. As shown in the results shown in Table 4, decreasing the L1/G1 ratio improves the overall performance of the operation, with less energy per ton of CO2 is utilized to carry out the electrolysis. Moreover, the amount of sea water utilized also decreases by decreasing the L1/G1 ratio, which will reduce the associated pumping costs (both capital and operational).

[0085] As the L1/G1 ratio decreases, the amount of CO2 mineralized (i.e.
converted into calcium or magnesium carbonates) decreases as a fraction of the total amount of CO2 captured. This is because sea water is originally saturated in CO2 and carbonates, and changing the final concentration of calcium and magnesium ions (through precipitation) shifts the saturation levels of carbonated ions and the solubility of CO2. At higher L1/G1 ratios (e.g., >70), more CO2 is precipitated as carbonate than what has been effectively removed from the flue gas, indicating that part of the CO2 originally in the sea water was co-precipitated.
At lower L1/G1 ratios (e.g., <60) more CO2 is removed from the flue gas than what is precipitated as solid carbonates, which indicates that the sea water has been enriched in carbonate ions and solubilized CO2. Consequently, it may be expected that part of this excess CO2 will came out of solution if the pH is lowered or if there are other processes that modify the final equilibrium points. Furthermore, as the spent sea water may be returned to the sea, it is desirable to return it at the same pH as it was taken up in order to inhibit the disturbance to the natural processes occurring in the sea.

[0086] To calculate the energy utilized for pumping sea water is necessary to first determine the total pressure drop in the liquid circuit (sea water pipeline and process equipment). For the purpose of a preliminary estimation, a 5 bar maximum pressure drop was assumed. As a result, the total energy cost for both pumping and electrolysis may be calculated per ton of CO2. The total cost per ton of captured, as well as per ton of CO2 mineralized is shown in the chart 400 shown in FIG. 4 (For a 0.0 V overvoltage).

[0087] The total CO2 captured is a contentious quantity, as it may be defined as either the total CO2 removed from the flue gas and stored in the sea water and in the solid carbonates, or it may be strictly defined as the CO2 that has been permanently sequestered (over geological time) in a mineral form. If the first definition is followed, then it may be concluded that the total energy consumption increases with increasing L/G ratio. However, if the more strict definition is followed, an optimum L/G ratio may be found (around 50, mass basis) at which the energy consumption per ton of CO2 captured is minimized. As pointed out previously, it is desirable to return the spent sea water to the sea at its original pH.
Consequently, the preferred L/G ratio will be slightly above the economic optimum point (between 50 and 70, mass basis).

[0088] The disclosed electrochemical CO2 capture process has chlorine gas as a by-product. Approximately 1.6 ton of Cl2 may be produced per ton of CO2 captured. Current worldwide chlorine production capacity is 60 million metric tonnes per year (see, for example, World Chlorine Council, 2012). Chlorine is mainly used for the production of polyvinyl chloride (35%), isocyanate and propylene oxide (15%), other organic derivatives (20%), inorganic chemicals (20%), and other uses including sanitation and water treatment (10%). The global market for PVC is expected to grow at a rate of 4-5 % per year.

[0089] Industrial production of chlorine is an energy intensive process that produces caustic soda (sodium hydroxide) as a by-product. In some areas, the produced soda does not have a local market and consequently must be transported over long distances, adding to the total cost. In these locations, the proposed process might be more attractive as only chlorine, but not caustic soda, is produced as a by-product of carbon capture.

[0090] The production of large amounts of chlorine associated with the electrochemical carbon capture system will nonetheless flood the current chlorine market. For instance, capturing only 0.12% of the global CO2 emissions (30.8 billion tons in 2009 (see, for example, Friedlingstein P., Houghton R.A., Marland G., Hackler J., Boden T.A., Conway T. J., Canadell J. G., Raupach M.R., Ciais P.
and Le Quere C. (2010) Update on CO2 emissions. Nature Geoscience 3, 811-812) would produce enough chlorine to supply all current worldwide consumption. Current Free on Board (FOB) prices for chlorine range from $800 to $1600 per ton and this price would likely collapse if the market were flooded. A reduction in chlorine prices, however, might increase the global demand for PVC and other chlorinated plastics.

[0091] In Table 5 included at Appendix 6 to this disclosure, a list of chlorine-containing polymers is presented together with the relative chlorine content in the monomer, and total global consumption. At 29924 kton/year production (Table 5), PVC is the main chlorinated polymer, with all other chlorinated polymers accounting for less than 3% of the chlorinated polymer market share.

[0092] Chlorine may also be used to produce hydrochloric acid than may later be used as a leaching agent in mineral extraction or simply be disposed of after neutralization with certain minerals. Chlorine gas reacts spontaneously with water in a disproportionation reaction:
Cl2 (g) + H2O HCI + HCIO ATP= -2.22 kJ/mol (24)

[0093] The hypochlorous acid is highly reactive and yields HCI in the presence of light or transition metal oxides of copper, nickel, or cobalt:
2 HCIO "V > 2 HCI + 02 AFP= 92.5 kJ/mol (25)

[0094] The hydrochloric acid may then be used to dissolve mineral deposits.
Of particular interest are the deposits of silicates that may be used as end donors of calcium and magnesium ions.

[0095] In the above disclosure, the source of calcium and magnesium cations was assumed to be from sea water; these cations could also be derived from land based rocks. Mineral carbonation is a natural process in which calcium and magnesium silicates react with atmospheric CO2 or CO2 dissolved in ground water to yield calcium or magnesium carbonate and silica (S102). Unfortunately, this natural process is very slow and only of significance over geological time frames.
Mining, grinding, and dispersion of silicates have been proposed as ways to accelerate mineral carbonation. These operations, however, may consume significant amounts of energy. Leaching of silicates with HCI has the potential to reduce energy utilization.

[0096] Hydrochloric acid, working in a closed loop manner, may be used to accelerate the in-situ underground conversion of silicates into silica and the produced calcium or magnesium brine may then be used for the production of calcium or magnesium carbonates at the surface. In this arrangement, fresh sea water consumption may be reduced to only make up the chloride losses to the underground silicate formation.

[0097] The reaction of silicates with HCI is in all cases exothermic and spontaneous. As an example, consider the following three silicates:
Mg2SiO4(S) + 4 HCI 2 MgCl2 SiO2() 2 H20 AH
¨236.9 kJ/mot (26) Ca2SiO4() + 4 HCI -4 2 CaC12 (aq) S102 (s) 2 H20 Ae= -255.3kYmol (27) MgSiO3 + 2 HCI MgC12 (am + Si02 (s) + H20 Ae= -109.3 k.1/mol (28)

[0098] Because the silicate dissolution reaction will occur underground, it is unlikely that the heat of reaction could be harvested. On the other hand, the heat generated by the chlorine dissolution reaction may potentially be recovered to at least partially offset the power consumption during electrolysis.

[0099] The global reaction for CO2 sequestration may be summarized as follows:
Ca2+ + CO2 (g) H2O CaCO3 + 2 H+
(29) 4 Fl+ + 02 (g) 4 a- 2 Cl2 + 2 H2 (30) 2 Cl2 (9) + 2 H20 ¨> 4 HCI + 02 (31) Ca2+ + CO2 (g) + H20 + 2 a- CaCO3 + 2 HCI Ae= -378.3 k.1/mol (32)

[00100] As shown above, the global reaction is exothermic; therefore, there is the potential (through efficiency improvements and thermal integration) to carry out the process without external energy sources. A schematic of a system 500 illustrating for this alternative process is shown in FIG. 5.

[00101] The system includes a CO2 absorption column 502, similar to CO2 absorption column 102 shown in FIG. 1, a electrochemical cell 504, similar to electrochemical cell 300 shown in FIG. 3, a separator 506, similar to separator 118 shown in FIG. 1, and a tank 508 the chlorine produced in the electrochemical cell 504 reacts to produce hydrochloric acid. Heat from the reaction occurring within the tank 508 may be recovered by a coil 510. Hydrochloric acid is deposited in a silicates deposit 512 to accelerate the in-situ underground conversion of silicates into silica and the produced calcium or magnesium brine may then be used for the production of calcium or magnesium carbonates at the surface, facilitating the operation of a closed loop to reduce fresh sea water consumption of the process.

[00102] As shown in reactions (26-28), for each mole of silicate that reacts, one mole of silica (Si02) is produced. The molar volume of silica is 22.68 cm3/mol, which is slightly more than 50% the molar volume of the magnesium silicate (43.02 cm3/mol) or calcium silicate (40.05 cm3/mol). Therefore, it is expected that as the leaching progresses, a cavern will be formed. Final abandonment of the cavern may be accomplished by flooding the cavern with carbonate slurry.

[00103] Deposits of magnesium and calcium silicate are relatively common.
In Canada, biotite and amphibolite -silicate rocks rich in magnesium and calcium-are a common occurrence in the bedrock of North Eastern Alberta, South and Central British Columbia, and Northern Saskatchewan. Higher quality deposits of wollastonite (CaSiO3) and olivine (Mg2S104) have also been reported in British Columbia and Quebec. Some of these deposits are currently being commercially exploited and others are reported as prospects. The Zippa Mountain (56039'10"N-131018'07"W) in British Columbia has estimated reserves totalling 20 Million tons of wollastonite, which could provide enough calcium to permanently sequester up to 7.58 million tons of CO2.

[00104] As described above, ammonia (NH3) may be used to assist the precipitation of calcium and magnesium carbonates. The resulting solution will be enriched in solubilised ammonium chloride (NH4CI). It is possible to regenerate the ammonia from this ammonium chloride solution, by passing the solution through an activated carbon (AC) column (see, for example, Huang, H.P. et al (2001) Dual alkali approaches for the capture and separation of CO2. Energy and Fuels 15(2), 263-268, (hereinafter "Huang")). The absorption of chloride and regeneration of ammonia is described as follows:
NH4CI + AC AC=FICI + NH3 AHads,õption 10.6 kJ/mol (33)

[00105] The activated carbon may later be regenerated by flowing steam (or hot water) through the activated carbon (see, for example, Huang):
AC.1-1C1 + H20 AC + HCI + H20 A I I desorption 24.8 kJ/mol (34)

[00106] The system 500 that carries out the carbon mineralization process previously described with reference to FIG. 5 may be modified to incorporate the ammonia-assisted carbonated precipitation and ammonia regeneration, as shown in the FIG. 6.

[00107] FIG. 6 shows a system 600 for CO2 mineralization with ammonia and ammonia regeneration. The system 600 includes a CO2 absorption column 602, similar to CO2 absorption column 502 shown in FIG. 5, a precipitation column in which carbonate precipitates with ammonia as discussed above, and ammonia regeneration tanks 606 and 608. Hydrochloric acid that exits the ammonia regeneration tanks 606 and 608 is deposited in a silicates deposit 610, similar to the deposition in the silicates deposit 512 described above with reference to FIG. 5.
Ammonia regeneration systems are common in several industries, and therefore the system 600 is not described in further detail herein.

[00108] The regeneration of ammonia described in disclosure is based on the results reported by Huang using activated carbon. It is possible, however, that other processes for regeneration may be more economical.

[00109] The carbonate precipitation with ammonia is an exothermic process, as has been shown in Eqs. (31-33). Based on the material balance equations presented before, the following energies may calculated:
= 1381 MJ will be released per ton of CO2 that is precipitated;
= 45.4 kmol of NH4 + should be regenerated per ton of CO2;
= 481 MJ per ton of CO2 will be released during ammonia regeneration;
= 1126 MJ per ton of CO2 are utilized for regeneration of the activated carbon;
= Total energy in: 1126 MJ/ton CO2; and = Total energy released: 1862 MJ/ton CO2.

[00110] Because more energy may be released during the process than the amount of energy utilized for regeneration of the activated carbon it is at least theoretically possible to design a system that will not require any thermal energy input. The challenge is to achieve a satisfactory energy integration scheme to allow the recovery of most of the energy released in the precipitation column and in the ammonia regeneration column.

[00111] A preliminary economic evaluation of the four alternative processes described above were carried out in terms of expected revenues and capital and operating costs. It was found that the four alternative processes described above have the potential to be economically feasible for permanently capturing CO2 emissions directly from flue gas.

[00112] The use of a catalyst is a desirable to reduce the size of the CO2 absorption equipment (and associated capital costs) as well as to reduce total operating costs. Catalysts must be selected based on cost, stability, and activity level. The carbonic anhydrase enzyme is the most efficient catalyst known for the CO2 hydration reaction in terms of activity but the stability of the enzyme under industrial operating conditions and the associated costs of the enzyme are unknown. As an alternative, nickel nanoparticles have been shown capable of catalyzing the hydration reaction, but the activity of the nanoparticles was several times lower than that of the carbonic anhydrase enzyme (see, for example, Ouled Ameur, Z. and Husein, M. (2013) Electrochemical behavior of potassium ferricyanide in aqueous and (w/o) microemulsion in the presence of dispersed nickel nanopartciles. Separation Science and Technology 48(5), 681 689).
Therefore, it may be desirable to conduct laboratory and pilot tests, using sea water or brines and flue-gas to evaluate the performance of several potential catalysts to determine the most cost effective one.

[00113] Permanent sequestration of carbon dioxide as solid carbonates may be a feasible solution to the increased levels of CO2 in the atmosphere.
Accelerated carbonation may be achieved through bio-mimicking approaches in which catalysts are used to accelerate the dissolution of CO2 in water and the formation of bicarbonate ions, while membrane or chemical systems may be used to remove the excess hydronium ions that are formed and that impede carbonate precipitation.

The economic feasibility of electrochemically assisted precipitation of carbonates may be increased if the process is coupled to the production and commercialization of chlorine, hydrochloric acid, or polyvinyl chloride.

[00114] Sequestering solid carbonates may involve depositing solid carbonates in shallow waters to create artificial islands or expand shore lines (similar to the effect of coral reefs). Carbonate brines may be deposited in caverns or other underground formations or utilized to create new hills or mountains. Another possible option is to utilize these CO2 capture products of as construction materials.

[00115] The described embodiments are to be considered in all respects only as illustrative and not restrictive. The scope of the claims should not be limited by the preferred embodiments set forth in the examples, but should be given the broadest interpretation consistent with the description as a whole. All changes that come with meaning and range of equivalency of the claims are to be embraced within their scope.

Appendix 1 Table 1 ¨ Typical composition of natural sea water, used for simulating the equilibrium condition for the CO2 ¨ sea water system Ion Mol/kg Cl- 0.54586 S042- 0.02824 Br- 0.00084 F- 0.00007 Na + 0.46906 0.05282 Ca2+ 0.01028 K+ 0.01021 Sr2+ 0.00009 HCO3- 0.00177 C032- 0.00026 B(0H)3 0.00032 B(OH)4- 0.00010 OH- 0.00001 CO2(g) 0.00001 Appendix 2 Geochemical Modelling Based - PHREEQC 3.0 output for equilibrium between sea water and flue gas. A. Carbonate saturation levels for original sea water; B.
Equilibrium conditions after contact with flue gas; C. Equilibrium conditions after neutralization with NaOH.
A. Initial solution 1. Sea water pH = 8.220 pe = 4.000 Total CO2 (mol/kg) = 2.131e-03 Temperature (deg C) = 25.000 Electrical balance (eq) = 1.308e-13 Phase SI log IAP log KT
Aragonite 0.59 -7.74 -8.34 CaCO3 Calcite 0.74 -7.74 -8.48 CaCO3 CO2(g) -3.39 -4.86 -1.47 CO2 Dolomite 2.36 -14.73 -17.09 CaMg(CO3)2 Strontianite -0.53 -9.81 -9.27 SrCO3 B. Reaction step 1. Using solution 1.
Solution after equilibrium with gas phase 1. Flue-gas pH = 5.838 Charge balance pe = 14.948 Adjusted to redox equilibrium Total CO2 (mol/kg) = 6.045e-03 Temperature (deg C) = 25.000 Electrical balance (eq) = 1.308e-13 Phase SI log IAP log KT
Aragonite -1.69 -10.02 -8.34 CaCO3 Calcite -1.54 -10.02 -8.48 CaCO3 CO2(g) -0.91 -2.37 -1.47 CO2 Dolomite -2.20 -19.29 -17.09 CaMg(CO3)2 Strontianite -2.82 -12.09 -9.27 SrCO3 C. Mixture 1: Neutralization of acidified sea water 9.950e-01 Solution 1 Solution after simulation 2.
5.000e-03 Solution 2 Caustic pH = 8.302 Charge balance pe = 12.484 Adjusted to redox equilibrium Total CO2 (mol/kg) = 6.015e-03 Temperature (deg C) = 25.000 Electrical balance (eq) = -1.887e-03 Phase SI log IAP log KT
Aragonite 1.11 -7.22 -8.34 CaCO3 Calcite 1.26 -7.22 -8.48 CaCO3 CO2(g) -3.03 -4.49 -1.47 CO2 Dolomite 3.40 -13.69 -17.09 CaMg(CO3)2 Strontianite -0.02 -9.29 -9.27 SrCO3 Appendix 3 Table 2 - Required reactor volumes for a single stage as a function of CO2 conversion.
Equilibrium conversion = 0.59. Input flow rate = 1 MM ton of CO2 per year.
Conversion CSTR Reactor Volume (m3) PFR Reactor Volume (m3) Enzyme Nickel Nickel Base Base Enzyme % of Xeci . (70- nanop. nanop.
case case (70-fold) fold) (14-fold) (14-fold) 90% , 24732 353 1767 6538 93 467 80% 11104 159 793 4623 66 330 70% 6544 93 467 3484 50 249 60% 4250 61 304 2671 38 191 50% 2863 41 205 2034 29 145 ¨
40% 1929 28 138 1509 22 108 30% 1253 18 90 1060 15 76 20% 739 11 53 667 10 48 Appendix 4 Table 3 - Inlet and outlet concentrations for a sequestration system with pH
regulation and L/G = 1000 Sea Spent Variable Units water Flue Sea Treated gas gas feed water Flowrate kg/[t] 1 0.00100 0.999408 0.00087 6.31E-[H+] mol/kg 09 1.28E-08 1.12E-[CO2*] mol/kg 05 2.943 1.12E-05 1.3E-05 2.59E-[HCO3-] mol/kg 03 1.28E-03 4.61E-[C032] mol/kg 04 1.13E-04 [Ca2+] mol/kg 0.01028 0.00569 [Mg2+] mol/kg 0.05282 5.29E-02 [Na] mol/kg 0.46906 0.46933 [Cr] mol/kg 0.54586 0.53901 [H20] mol/kg 53.5559 53.585 9.61E-[Ohl-] mol/kg 06 4.75E-06 pH 8.2 7.89 [CO2(g)] %molar 0.0394% 8.210% 0.0394%
CaCO3 (s) = 0.00459 mol/[t]
H+ = 7.165x10-3 mol/[t]

Appendix 5 Table 4 - Power utilization for electrolysis and sea water utilization for the electrolytic precipitation of carbonates at different UG ratios 0.8 V
0 V Overvolt. Sea Electro- Overvolt. CO2 water Final ns Net P
mineraliz Lo/GoNet P Elec. Elec. flowrat pH utilized (0.8 ed (OV) Cost V) Cost kmol GJ GJ m3 1000 7.89 55.6 0.7 12.15 5.0 83.64 7756 156.6%
500 7.93 50.1 0.7 10.95 4.5 75.38 3878 126.1%
400 7.94 49.1 0.6 10.73 4.4 73.90 3102 120.6%
300 7.95 48.1 0.6 10.52 4.3 72.42 2327 115.1%
200 7.97 47.1 0.62 10.3 4.3 70.91 1551 109.5%
100 8.04 46.1 0.60 10.1 4.2 69.31 776 103.5%
80 8.10 45.8 0.60 10.0 4.1 68.92 620 102.0%
70 8.14 45.6 0.60 10.0 4.1 68.67 543 101.1%
60 8.22 45.4 0.6 9.92 4.1 68.31 465 99.7%
50 8.39 44.9 0.6 9.81 4.1 67.57 388 96.5%
40 8.74 43.5 0.6 9.50 3.9 65.43 310 84.4%
25 9.14 41.6 0.5 9.10 3.8 62.62 194 53.7%
20 9.25 41.4 0.5 9.05 3.7 62.29 155 43.0%
9.50 41.7 0.5 9.11 3.8 62.75 78 21.5%
Note 1: All values (except pH) are given per ton of CO2 captured (in the sea water and as solid carbonates).

Appendix 6 Table 5 - List of chlorine-containing polymers and the relative chlorine content in the monomer and total global consumption.
Polymer Chlorine content World Production (mass %) kton / year Year PDCA 74.7% -O
PVDC 73.1% 160 2004 CPVC 56 - 74%
PVC 56.7% 29924 2009 Neoprene 40.0% 360 2012 CPE 25 - 40% 260 2003 CSPE 27 - 35%
PCTFE 30.4% 0.4 1998

Claims (20)

Claims

1. A method of separating carbon dioxide from a flue gas, the method comprising:
exposing a flue gas comprising carbon dioxide to salt water to generate a solution;
monitoring pH of the solution;
when a pH condition is met, adjusting the pH to precipitate a solid carbonate material from the solution;
separating the solid carbonate material from the solution; and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

2. The method according to claim 1, wherein the salt water is sea water.

3. The method according to claim 1, comprising adding a weak base to the salt water to form a buffered solution prior to exposing the flue gas to the salt water.

4. The method according to claim 1, wherein the pH condition is a reduction in pH.

5. The method according to claim 1, wherein adjusting the pH comprises increasing the pH by adding a base.

6. The method according to claim 1, wherein adjusting the pH comprises increasing the pH utilizing an electrochemical process.

7. The method according to claim 6, wherein the electrochemical process comprises removing protons utilizing a membrane.

8. The method according to claim 1, wherein the solid carbonate material comprises calcite, aragonite, dolomite, magnesite, or a combination thereof.

9. The method according to claim 6, wherein the ratio of sea water to flue gas is selected to minimize the energy utilized for the electrochemical process.

10. The method according to claim 6, wherein chlorine or vinyl chloride monomer are produced and used as starting materials for synthesizing other chemical products such as hydrochloric acid or polyvinyl chloride.

11. The method according to claim 1, comprising a catalyst for accelerating dissolution of the carbon dioxide in the salt water.

12. A method of separating carbon dioxide from a flue gas, the method comprising:
exposing an underground silicate deposit to hydrochloric acid to generate a brine;
recovering the brine at surface;
exposing the flue gas comprising carbon dioxide to the brine;
monitoring pH of the brine;
when a pH condition is met, adjusting the pH;
precipitating a solid carbonate material from the brine;
separating the solid carbonate material from the brine; and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

13. The method according to claim 12, comprising adding salt water to the brine before exposing the flue gas to the brine.

14. The method according to claim 12, wherein adjusting the pH comprises increasing the pH utilizing an electrochemical process.

15. The method according to claim 14, comprising producing chlorine from the electrochemical process and re-generating hydrochloric acid.

16. The method according to claim 12, wherein storing the solid carbonate material comprises delivering the solid carbonate material to the underground silicate deposit.

17. A method of separating carbon dioxide from a flue gas, the method comprising:
exposing an underground silicate deposit to hydrochloric acid to generate a brine;
recovering the brine at surface;
exposing the flue gas comprising carbon dioxide to the brine;
introducing ammonia to the brine to generate an ammonium salt and a solid carbonate material;
separating the solid carbonate material from the brine; and storing the solid carbonate material to inhibit the carbon dioxide from entering the atmosphere.

18. The method according to claim 17, comprising regenerating the ammonia from the ammonium salt.

19. A method of separating carbon dioxide from a flue gas, the method comprising:
exposing the flue gas comprising carbon dioxide to salt water to provide a solution;
monitoring pH of the solution;
adjusting the pH to generate a basic solution and precipitate a solid carbonate material from the solution;
separating the solid carbonate material from the solution; and storing the solid carbonate material to inhibit carbon dioxide from entering the atmosphere.

20. A method of sequestering carbon dioxide, the method comprising:
exposing a flue gas comprising carbon dioxide to sea water to provide a solution;
monitoring the pH of the solution;
when a pH condition is met, adjusting the pH to precipitate a solid carbonate material from the solution;
storing the solid carbonate material in the sea water to inhibit carbon dioxide from entering the atmosphere.