ZeroCAL: Eliminating Carbon Dioxide Emissions from Limestone's Decomposition to Decarbonize Cement Production

Abstract

Limestone (calcite, CaCO₃) is an abundant and cost-effective source of calcium oxide (CaO) for cement and lime production. However, the thermochemical decomposition of limestone (∼800 °C, 1 bar) to produce lime (CaO) results in substantial carbon dioxide (CO_2(g)) emissions and energy use, i.e., ∼1 tonne [t] of CO₂ and ∼1.4 MWh per t of CaO produced. Here, we describe a new pathway to use CaCO₃ as a Ca source to make hydrated lime (portlandite, Ca(OH)₂) at ambient conditions (p, T)-while nearly eliminating process CO_2(g) emissions (as low as 1.5 mol. % of the CO₂ in the precursor CaCO₃, equivalent to 9 kg of CO_2(g) per t of Ca(OH)₂)-within an aqueous flow-electrolysis/pH-swing process that coproduces hydrogen (H_2(g)) and oxygen (O_2(g)). Because Ca(OH)₂ is a zero-carbon precursor for cement and lime production, this approach represents a significant advancement in the production of zero-carbon cement. The Zero CArbon Lime (ZeroCAL) process includes dissolution, separation/recovery, and electrolysis stages according to the following steps: (Step 1) chelator (e.g., ethylenediaminetetraacetic acid, EDTA)-promoted dissolution of CaCO₃ and complexation of Ca²⁺ under basic (>pH 9) conditions, (Step 2a) Ca enrichment and separation using nanofiltration (NF), which allows separation of the Ca-EDTA complex from the accompanying bicarbonate (HCO₃ ^-) species, (Step 2b) acidity-promoted decomplexation of Ca from EDTA, which allows near-complete chelator recovery and the formation of a Ca-enriched stream, and (Step 3) rapid precipitation of Ca(OH)₂ from the Ca-enriched stream using electrolytically produced alkalinity. These reactions can be conducted in a seawater matrix yielding coproducts including hydrochloric acid (HCl) and sodium bicarbonate (NaHCO₃), resulting from electrolysis and limestone dissolution, respectively. Careful analysis of the reaction stoichiometries and energy balances indicates that approximately 1.35 t of CaCO₃, 1.09 t of water, 0.79 t of sodium chloride (NaCl), and ∼2 MWh of electrical energy are required to produce 1 t of Ca(OH)₂, with significant opportunity for process intensification. This approach has major implications for decarbonizing cement production within a paradigm that emphasizes the use of existing cement plants and electrification of industrial operations, while also creating approaches for alkalinity production that enable cost-effective and scalable CO₂ mineralization via Ca(OH)₂ carbonation.

Figure 1

Speciation, in aqueous solution at 25 °C and 1 bar, of (a) CaCO₃ and Ca(OH)₂ and (b) Dissolved Inorganic Carbon (DIC) and EDTA. For reference, the limiting solubilities of CaCO₃, CO₂, and EDTA in water at pH 7 are ∼4.0, ∼0.1, and ∼420 mmol/L, respectively.

Figure 2

Representative process flow diagram (PFD) of the ZeroCAL process showing the different unit operations and materials flows.

Figure 3

Dissolution of calcite at room temperature in the presence of EDTA. (a) Increased solubility of calcite (i.e., Ca concentration) as a function of the pH across a range of EDTA concentrations. The black dashed line shows the thermodynamically modeled solubility of calcite in a system devoid of EDTA. (b) Degassing of CO₂ expressed as a molar percentage of total amount of aqueous and mineralized carbonates in the system as a function of the EDTA concentration and the initial pH of the solution. The dashed black line shows the thermodynamically modeled CO₂ emission from calcite in a system devoid of EDTA, and the blue dotted line shows the CO₂ emission of calcite in a system containing EDTA calculated using eqs 12–14. (c) Ca concentration at equilibrium at pH 9.5 and in the presence of 100 mmol/L EDTA in MQW, a 0.5 mol/L NaCl solution, and simulated seawater (“Instant Ocean Seawater: IOSw”) for reagent calcite (AR calcite) or limestone rock.

Figure 4

(a) “Initial” dissolution rate of calcite at high undersaturation (i.e., far from dissolution equilibrium) as a function of the initial pH of the solution across a range of EDTA concentrations. The data for dissolution occurring in MQW (“0 mmol/L EDTA”) are taken from the literature.^,, (b) Measured and fitted (eq 3) relationship between the dissolution rate and the [Ca]/[EDTA] (molar) ratio. (c) Dissolution rates of reagent-grade calcite and limestone rock at pH 9.5 and in the presence of 100 mmol/L EDTA in MQW, 0.5 mol/L NaCl, and simulated seawater (IOSw).

Figure 5

(a) Experimental performance of the selected NF membranes for Ca/CO₂ [Ca-EDTA and Na⁺ + HCO₃^–] separation at 0% water recovery. Here, R_Ca-EDTA and R_Na-HCO3 indicate the rejection of the Ca-EDTA^2– and Na⁺-HCO₃^– species, respectively. An optimal membrane maximizes R_Ca-EDTA, minimizes R_Na-HCO, and ensures a suitable flux at increasing water recovery. (b) Upon initial screening, Ca/CO₂ separation and flux of the remaining membrane candidates relative to the XN45 membrane up to 50% water recovery are shown, highlighting the Ca/CO₂ separation properties of XN45. (c) XN45 membrane’s separation of Ca/CO₂ at 85% water recovery for limestone rock and AR-grade calcite precursors at pH 8 and 10, for single- and double-stage separation, using water and a high salinity simulated seawater matrix.

Figure 6

(a) FT-IR spectra of the Ca-EDTA complex for pHs ranging from 2 to 7 representing the regions of stability (complexation) and instability (decomplexation). (b) EDTA recovery efficiency as a function of the initial Ca-EDTA concentrations and time of acid exposure. (c) Ca decomplexation and EDTA recovery efficiencies for different Ca-EDTA concentrations after 90 min. In general, for EDTA concentrations >100 mmol/L, near stoichiometric recovery of the EDTA is observed within 5 min.

Figure 7

Electrolyzer performance for a Faradaic efficiency of 90% and a cell voltage of −2.06 V showing (a) Ca conversion (into Ca(OH)₂) and (b) the electric energy intensity (EEI) of Ca(OH)₂ precipitation. Both parts a and b refer to Ca(OH)₂ precipitation (Step 3) only and are shown as a function of the inlet Ca concentration and the current/flow rate ratio. (c) Purity of portlandite produced as a function of the Ca/CO₂ separation in solution after Step 2a-i. (d) Extent of CO₂ degassing (as a percent of the amount of CO₂ introduced into the system via the dissolution of limestone) that could occur as a function of the initial EDTA concentration and extent of Ca/CO₂ separation after Step 2a-i.

Figure 8

(a) CO₂ evolved (process emissions) across different configurations. (b) Process water demand, (c) gross energy intensity, and (d) net energy intensity of the ZeroCAL process considering Steps 0–3 across three different configurations (C1, C2, and C3) that encompass single- or double-stage NF (i.e., to minimize CO₂ evolution), with or without process water recycling (i.e., to reduce the water demand and/or the non-electrolysis process energy) for an initial Ca-EDTA concentration of 100 mmol/L. Herein, in part c, electrolytic Ca(OH)₂ precipitation (Step 3) is estimated to have a gross energy intensity of ∼1.7 MWh/t_Ca(OH)2, while the remainder of the energy is attributed to “upstream” operations (Steps 0–2).

Figure A1

Schematic representation of the reactions and species migrations occurring in the electrolyzer. Species migrations are specified for the different types of membranes/separators that may be used to separate the anolyte and the catholyte, e.g., porous membrane (non-ion specific), cation exchange membrane (CEM), and anion exchange membrane (AEM).

Figure A2

Schematic representation of the chemical boundary conditions considered in the flow electrolyzer.

Figure A3

(a) Dissolution rate of calcite as a function of the solution’s ionic strength for a starting pH of 9.5. (b) Performance of the recovered EDTA in chelating Ca extracted from calcite (CaCO₃) compared to “pristine” AR-grade EDTA over 5 cycles of use and recovery.

Figure A4

Performance of the XN45 membrane on single-stage NF with a pH 8 Ca-EDTA/Na-HCO₃ solution system as a function of water recovery comparing AR-grade calcite and natural limestone feedstocks: (a) Ca-EDTA and Na-HCO₃ real rejections; (b) profile of the cumulative HCO₃ permeation; (c) normalized permeate flux; (d) pH in the streams.

Figure A5

Effect of discharging a pH 9 stream containing 100 mmol/L aqueous Na-HCO₃ into seawater as a function of the considered dilution factor (the lines are thermodynamically modeled, and the points were acquired experimentally): (a) the carbonate evolution (as CO₂), (b) the concentration of Mg, Ca, and the carbonate species, and (c) the amount of Ca and Mg precipitated, highlighting the formation of calcium carbonate and the absence of hydrated magnesium carbonate formation in low dilution condition. The arrow indicates that the model fails to accurately represent the natural supersaturated state of seawater.

Figure A6

Performance of the NF90 NF membrane in Step 2a-ii on a Na-HCO₃-containing solution from the previous separation Step 2a-i: (a) instantaneous Na-HCO₃ real rejection and concentration factor (CF) in the retentate stream; (b) pH in the streams.

Figure A7

(a) XRD and FT-IR normalized patterns of the recovered EDTA precipitate and AR-grade H₄EDTA. (b) Conservative estimate of the dissolved carbonate degassed from solution (as CO₂) as a function of the initial carbonate concentration during the pH-swing from pH 8–10 to pH 2. No EDTA is present in the system, which removes the (low) dilution effect induced by EDTA titration.

Figure A8

(a) Comparison between the experimental and modeled portlandite purity at ∼100% Ca conversion for different Ca/CO₂ separation values. (b) Contaminants (Al, Fe, Si, and Mg) removed from the Ca-enriched feed by precipitation of hydroxides and hydrated solids. SiO₂ is amorphous silica, and C-M-S-H is a calcium and magnesium containing silicate hydrate phase.