2023 Volume 64 Issue 12 Pages 2764-2775
A thermodynamic model (U-Cal model) is proposed to predict the fugacity and water content in the CO2-rich phase (gas or supercritical fluid), and the CO2 solubility and pH in aqueous solutions in a CO2 capture and storage environment, i.e., a CO2 environment in a supercritical state. The values predicted by the U-Cal model agree well with measured values. The water content in the CO2-rich phase increases to 1–10 g/L as a result of the mutual dissolution of CO2 and H2O. The increase in the solubility of CO2 in aqueous solution and the decrease in the pH with increasing pressure are small. The CO2 corrosion behavior of carbon steel is discussed to use the U-Cal model. In iron dissolution-dominant CO2 corrosion in an aqueous solution with carbon steel, the corrosion rate can be understood as a function of the pH. In FeCO3 formation-dominant CO2 corrosion, it is considered that the corrosion progresses as FeCO3 dissolves to supersaturation and then FeCO3 precipitates on the surface of the material. The FeCO3 precipitation behavior is predicted from the crystal growth rate equation. Corrosion of carbon steel in the CO2-rich phase involves similar mechanisms to corrosion in an aqueous solution; however, the corrosion rate is lower.
This Paper was Originally Published in Japanese in Zairyo-to-Kankyo 72 (2023) 131–141.
Efforts to stop the increase in carbon dioxide (CO2) emissions or reduce them to virtually zero will be required in the second half of this century as measures against global warming.1) CO2 capture and storage (CCS) technology is attracting attention as the most effective way to stop increases of CO2.2) In Japan, a large-scale CCS demonstration test was conducted in Tomakomai starting in April 2016, with the initial target of 300,000 tons of CO2 injection being achieved in November 2019; the development of CCS technology has been studied and related-technology data have been accumulated.3) It is well known in the fields of oil and gas production and transportation that wet CO2 environments are highly corrosive for carbon steels. Researches on CO2 corrosion began in the 1940s, this topic was actively studied in the 1980s and the studied-papers were compiled as a collection of papers by the National Association of Corrosion Engineers (NACE, now AMPP).4,5) The main features are acceleration of the cathodic reaction by supplying hydrogen ions from carbonic acid dissolved in the aqueous solution, and the anodic reaction by the dissolution of iron and formation of iron carbonate (FeCO3), which is a corrosion product. In addition to general corrosion, localized corrosion with FeCO3 formation (ringworm corrosion, mesa corrosion) is also a problem as a corrosion phenomenon. In CCS, CO2 is transported and injected underground under a supercritical state (the critical temperature of CO2 is 304.25 K, and the critical pressure is 7.37 MPa); this enables the CO2 to be handled economically and in large quantities. The CO2 corrosion which has been studied in conventional oil well fields occurs in an environment that can be approximated even if the CO2 gas and aqueous solution are treated as ideal conditions. Deep wells and enhanced oil recovery (EOR) are also subject to supercritical CO2 environments.6,7) To study supercritical CO2 environments, it is necessary to understand the behavior of real gases and real solutions.
In this paper, we combine the Peng-Robinson equation of state, which considers the mutual dissolution of two components, and is usually used in high-pressure environments, the pressure correction equation for Henry’s low constant of CO2, and an equation for estimating the activity coefficient in an aqueous solution to use the Pitzer model. A simple computational model is developed to use Microsoft Excel to predict the thermodynamic properties related to the corrosion environments in CCS. Using this computational model, the CO2 corrosion behavior in a supercritical CO2 environment and the precipitation of FeCO3 are investigated from the viewpoints of equilibrium theory and reaction theory.
A thermodynamic calculation model (U-Cal model) has been developed to predict the fugacity and H2O mole fraction in the CO2-rich phase (gas or supercritical fluid phase; Naq), and the solubility (mol/kg) of CO2 and the pH in the aqueous solution (Aq) phase as environmental factors. The equilibrium between the Naq and Aq phases is expressed as follows.
| \begin{equation} \text{CO$_{2}$(Naq phase)} \to \text{CO$_{2}$(Aq phase)},\quad \text{$K$(CO$_{2}$)} = \frac{\text{$y$(CO$_{2}$)}}{\text{$x$(CO$_{2}$)}} \end{equation} | (1) |
| \begin{equation} \text{H$_{2}$O(Naq phase)} \to \text{H$_{2}$O(Aq phase)}, \quad \text{$K$(H$_{2}$O)} = \frac{\text{$y$(H$_{2}$O)}}{\text{$x$(H$_{2}$O)}} \end{equation} | (2) |
Here, y(CO2) and y(H2O) are the mole fractions of CO2 and H2O, respectively, in the Naq phase; and x(CO2) and x(H2O) are the mole fractions of CO2 and H2O, respectively, in the Aq phase. The following relations hold in a NaCl solution:
| \begin{equation} \text{Naq phase:}\ \text{$y$(H$_{2}$O)} + \text{$y$(CO$_{2}$)} = 1 \end{equation} | (3) |
| \begin{equation} \text{Aq phase:}\ \text{$x$(H$_{2}$O)} + \text{$x$(CO$_{2}$)} + \Sigma x_{i} = 1 \end{equation} | (4) |
| \begin{equation} \Sigma x_{i} = \frac{\Sigma m_{i}}{(M_{\text{w}} + \Sigma m_{i} + \text{$m$(CO$_{2}$)})} \end{equation} | (5) |
The mole fraction is usually used in the Naq phase; accordingly, the mole fraction was also used in the Aq phase. mi is the molality (mol/kg) of the solute and Mw is the mole number in 1 kg/L of H2O, 55.508 mol. The amount of dissolved CO2, m(CO2), was assumed to be zero. The thermodynamic equilibrium of the Naq phase is expressed as
| \begin{equation} \mu_{\text{Naq}}(T,P) = \mu_{\text{Naq}}^{0}(T,P) + RT\ln(f),\quad f = P\varPhi y. \end{equation} | (6) |
Here, μ is the chemical potential, μ0 is the chemical potential in the standard state, R is the gas constant, T is the temperature in K, P is the total pressure in MPa and f is the fugacity in MPa. f is calculated from the mole fraction y of the Naq phase, the fugacity coefficient Φ, and the total pressure P. The chemical potential of the Aq phase is expressed as
| \begin{equation} \mu_{\text{Aq}}(T,P) = \mu_{\text{Aq}}^{0} (T,P) + RT\ln (a), \end{equation} | (7) |
where a is the activity. The chemical potentials of the Naq and Aq phases are equal in the equilibrium state.
| \begin{equation} \frac{\mu_{\text{Naq}}^{0}(T,P) - \mu_{\text{Aq}}^{0}(T,P)}{RT} = \ln \left(\frac{a}{P\varPhi y} \right) \end{equation} | (8) |
Then, the following equations were derived for H2O and CO2:
| \begin{equation} K'(\text{H$_{2}$O})x(\text{H$_{2}$O}) = P\varPhi(\text{H$_{2}$O})y(\text{H$_{2}$O}), \end{equation} | (9) |
| \begin{equation} H(\text{CO$_{2}$})x(\text{CO$_{2}$})\gamma_{\text{x}}(\text{CO$_{2}$}) = P\varPhi(\text{CO$_{2}$})y(\text{CO$_{2}$}). \end{equation} | (10) |
Here, Φ(H2O) and Φ(CO2) are the fugacity coefficients of H2O and CO2, respectively, in the Naq phase; H(CO2) is Henry’s constant of the CO2 gas; K′(H2O) is a constant given by eq. (11) and γx(CO2) is the activity coefficient of the CO2 mole fraction in the aqueous solution. The activity of H2O was approximated as being equal to the H2O mole fraction. Using eq. (9), the equilibrium constant K(H2O) between the Naq and Aq phases with respect to H2O is expressed as
| \begin{align} K(\text{H$_{2}$O})& = \frac{y(\text{H$_{2}$O})}{x(\text{H$_{2}$O})} = \frac{K'(\text{H$_{2}$O})}{P\varPhi(\text{H$_{2}$O})} \\ &= \frac{K_{0}(\text{H$_{2}$O})\exp((P - 1)V(\text{H$_{2}$O})/RT)}{P\varPhi(\text{H$_{2}$O})}, \end{align} | (11) |
where V(H2O) is the average partial molar volume of H2O, which is 18.18 cm3/mol. K0(H2O) is the equilibrium constant at the reference pressure of 0.1 MPa, with the following equation proposed by Ganji et al.8) adopted here:
| \begin{align} \log(K_{0}(\text{H$_{2}$O})) &= -2.209 + 3.097 \times 10^{-2} \times \theta - 1.098\\ &\quad \times 10^{-4} \times \theta^{2} + 2.048 \times 10^{-7} \times \theta^{3}, \end{align} | (12) |
where θ is the temperature in degrees Celsius. From eqs. (2) and (4), the mole fraction of H2O in the Naq phase is obtained such that
| \begin{align} y(\text{H$_{2}$O}) &= x(\text{H$_{2}$O})K(\text{H$_{2}$O}) \\ &= (1 - x(\text{CO$_{2}$}) - \Sigma x_{i}) K(\text{H$_{2}$O}). \end{align} | (13) |
Using eq. (10), the equilibrium constant K(CO2) between the Naq and Aq phases with respect to CO2 is expressed as
| \begin{equation} \text{$K$(CO$_{2}$)} = \frac{\text{$y$(CO$_{2}$)}}{\text{$x$(CO$_{2}$)}} = \frac{\text{$H$(CO$_{2}$)}\gamma_{\text{x}}(\text{CO$_{2}$})}{P\varPhi(\text{CO$_{2}$})}. \end{equation} | (14) |
Here, H(CO2) is Henry’s constant and γx(CO2) is the activity coefficient for the molar fraction, which was converted from the activity coefficient for the molality γ(CO2) using the following relational expression in which m(CO2) was assumed to be zero:
| \begin{equation} \gamma_{\text{x}}(\text{CO$_{2}$}) = \frac{M_{\text{w}} + \Sigma m_{i} + \text{$m$(CO$_{2}$)}}{M_{\text{w}}}\gamma(\text{CO$_{2}$}). \end{equation} | (15) |
Using eqs. (1) and (3), the solubility of CO2 in the Aq phase is calculated in mole fraction such that
| \begin{equation} \text{$x$(CO$_{2}$)} = \frac{1 - \text{$y$(H$_{2}$O)}}{\text{$K$(CO$_{2}$)}}. \end{equation} | (16) |
Accordingly, the solubility of CO2 in molality (mol/kg), m(CO2), can be obtained by converting from x(CO2) such that
| \begin{equation} m(\text{CO$_{2}$}) = \frac{x(\text{CO$_{2}$})(\Sigma m_{i} + M_{\text{w}})}{1 - x(\text{CO$_{2}$})}. \end{equation} | (17) |
y(H2O) in the Naq phase can be obtained from eqs. (13) and (16):
| \begin{equation} y(\text{H$_{2}$O}) = \cfrac{1 - \cfrac{1}{K(\text{CO$_{2}$})} - \Sigma x_{i}}{\cfrac{1}{K(\text{H$_{2}$O})}-\cfrac{1}{K(\text{CO$_{2}$})}}. \end{equation} | (18) |
Then, the fugacity of H2O, f(H2O), in the CO2-rich phase (gas or supercritical fluid phase; Naq phase) is PΦ(H2O)y(H2O), and the fugacity of CO2, f(CO2), is PΦ(CO2)y(CO2). When considering corrosion, it is easier to understand the water content (grams of H2O in 1 L of the Naq phase); accordingly, this quantity is calculated from y(H2O) and the compression factor Z (= Pv/RT, where v is the volume of Naq phase).
pH is an important factor in corrosion in aqueous solutions. Therefore, in addition to the dissolution reaction of CO2, the pH was calculated to consider the following ionic equilibria and equilibrium constants.
| \begin{align} &\text{CO$_{2}$(aq)} + \text{H$_{2}$O} \to \text{H$^{+}$} + \text{HCO$_{3}{}^{-}$}\\ &\quad K_{10} = a(\text{H$^{+}$})a(\text{HCO$_{3}{}^{-}$})/(a(\text{CO$_{2}$(aq)})a(\text{H$_{2}$O})) \end{align} | (19) |
| \begin{align} K_{1} &= m(\text{H$^{+}$})m(\text{HCO$_{3}{}^{-}$})/m(\text{CO$_{2}$(aq)}) \\ &= K_{10}(a(\text{H$_{2}$O})\gamma(\text{CO$_{2}$})/(\gamma(\text{H$^{+}$})\gamma(\text{HCO$_{3}{}^{-}$}))) \end{align} | (20) |
| \begin{align} &\text{HCO$_{3}{}^{-}$} \to \text{H$^{+}$} + \text{CO$_{3}{}^{2-}$} \\ &\quad K_{20} = a(\text{H$^{+}$})a(\text{CO$_{3}{}^{2-}$})/a(\text{HCO$_{3}{}^{-}$}) \end{align} | (21) |
| \begin{align} K_{2}& = m(\text{H$^{+}$})m(\text{CO$_{3}{}^{2-}$})/m(\text{HCO$_{3}{}^{-}$}) \\ &= K_{20}(\gamma(\text{HCO$_{3}{}^{-}$})/(\gamma(\text{H$^{+}$})\gamma(\text{CO$_{3}{}^{2-}$}))) \end{align} | (22) |
| \begin{equation} \text{H$_{2}$O} \to \text{H$^{+}$} + \text{OH$^{-}$} \quad K_{\text{W0}} = a(\text{H$^{+}$})a(\text{OH$^{-}$})/a(\text{H$_{2}$O}) \end{equation} | (23) |
| \begin{equation} K_{\text{W}} = m(\text{H$^{+}$})m(\text{OH$^{-}$}) = K_{\text{W0}}(a(\text{H$_{2}$O})/(\gamma(\text{H$^{+}$})\gamma(\text{OH$^{-}$}))) \end{equation} | (24) |
| \begin{equation} \text{NaHCO$_{3}$} \to \text{Na$^{+}$} + \text{HCO$_{3}{}^{-}$} \end{equation} | (25) |
Here, m, a, and γ represent the molality, activity (a = mγ), and activity coefficient, respectively. HCO3− is assumed to be generated by dissolving NaHCO3 minerals. The following relationship exists among the measured values of m0(Na+) for Na+, m0(Cl−) for Cl−, and m0(NaHCO3) for HCO3−:
| \begin{equation} m_{\text{0alk}} = m_{0}(\text{NaHCO$_{3}$}) = m_{0}(\text{Na$^{+}$}) - m_{0}(\text{Cl$^{-}$}). \end{equation} | (26) |
Electrical neutrality for the total number of charges of cations and anions is given as
| \begin{align} m(\text{H$^{+}$}) + m_{0}(\text{Na$^{+}$}) &= m_{0}(\text{Cl$^{-}$}) + m(\text{HCO$_{3}{}^{-}$}) \\ &\quad + 2m(\text{CO$_{3}{}^{2-}$}) + m(\text{OH$^{-}$}). \end{align} | (27) |
Using the equilibrium constants in eqs. (29)–(32), eq. (27) is converted to the following eq. (28) for m(H+) only; m(H+) is then obtained by the iterative method. The pH is obtained by multiplying m(H+) by γ(H+) and taking the negative logarithm.
| \begin{align} \mathit{func}(m(\text{H$^{+}$})) &= m(\text{H$^{+}$})^{3} + m_{\text{0alk}}m (\text{H$^{+}$})^{2} \\ &\quad - (K_{\text{W}} + K_{\text{A}}f(\text{CO$_{2}$}))m(\text{H$^{+}$}) \\ &\quad - 2K_{\text{B}}f(\text{CO$_{2}$}) = 0 \end{align} | (28) |
| \begin{equation} K_{\text{A}} = K_{\text{H0}}K_{10}(a(\text{H$_{2}$O})/(\gamma(\text{H$^{+}$})\gamma(\text{HCO$_{3}{}^{-}$}))) \end{equation} | (29) |
| \begin{equation} K_{\text{B}} = K_{\text{H0}}K_{10}K_{20}(a(\text{H$_{2}$O})/(\gamma(\text{H$^{+}$})^{2}\gamma(\text{CO$_{3}{}^{2-}$}))) \end{equation} | (30) |
| \begin{equation} K_{\text{W}} = K_{\text{W0}}(a(\text{H$_{2}$O})/(\gamma(\text{H$^{+}$})\gamma(\text{OH$^{-}$}))) \end{equation} | (31) |
| \begin{equation} K_{\text{H0}} = M_{\text{W}}/H(\text{CO$_{2}$}) \end{equation} | (32) |
| \begin{equation} \text{pH} = -\log(\gamma(\text{H$^{+}$})m(\text{H$^{+}$})) \end{equation} | (33) |
The fugacity coefficient is calculated to use the Peng-Robinson equation of state for a binary mixture system extended under high-pressure conditions; this is an extension of the Van der Waals equation of state.9)
| \begin{equation} P = \frac{RT}{v - b} - \frac{a(T)}{v(v + b) + b(v - b)} \end{equation} | (34) |
This equation is rewritten in the form of a cubic equation of the compression factor Z (= Pv/RT) to obtain a numerical solution for Z. The fugacity coefficient Φ is defined in eq. (35) and the solution of the fugacity coefficient is obtained from eq. (36).
| \begin{equation} \ln(\varPhi) = \int_{0}^{P} \frac{Z - 1}{P}dP \end{equation} | (35) |
| \begin{align} \ln(\varPhi_{i})& = \frac{b_{i}}{b}(Z - 1)- \ln(Z - B) \\ &\quad - \frac{A}{2\sqrt{2}B}\left(\dfrac{2\displaystyle\sum\nolimits_{n = j}(y_{j}{a}_{ji})}{a} - \frac{b_{i}}{b} \right)\ln\left(\frac{Z + 2.414B}{Z - 0.414B} \right) \end{align} | (36) |
| \begin{equation} a_{ij} = \sqrt{a_{i}a_{j}}(1 - \delta_{ij}),\quad a = \Sigma \Sigma y_{i}y_{j}a_{ij},\quad b = \Sigma y_{i}b_{i} \end{equation} | (37) |
Here, A = aP/(RT)2, B = bP/RT, yi is the mole fraction in the Naq phase of H2O and CO2, ai (a of component i) is a coefficient related to intermolecular attractive interactions, bi (b of component i) is a coefficient related to repulsive interactions, and δij is the experimentally obtained interaction coefficient between the components i and j, which is 0.19014 according to Li et al. for the CO2 and H2O system.8)
2.2 Calculation of Henry’s law constant H(CO2)Henry’s low constant was corrected for temperature and pressure using the correction formula proposed by Akinfiev et al.:10)
| \begin{align} \ln(H\text{(CO$_{2}$)}) &= (1 - \eta) \ln (f_{\text{H2O}}^{0})\\ &\quad + \eta \ln \left(\frac{RT}{m_{\text{w(H2O)}}}\rho_{\text{H2O}}^{0} \right) + 2\rho_{\text{H2O}}^{0}\Delta B, \end{align} | (38) |
| \begin{equation} \Delta B = \tau + \beta \left(\frac{10^{3}}{T}\right)^{0.5}. \end{equation} | (39) |
Here, $f_{\text{H2O}}^{0}$ and $\rho_{\text{H2O}}^{0}$ are the fugacity and density, respectively, of pure water as functions of the temperature and pressure; η, τ, and β are correction factors; and mw(H2O) is the molar mass of water. These calculations are performed using the equations and coefficients used by Shabani et al.11)
2.3 Calculation of the equilibrium constants K10, K20 and KW0Using the relational expressions and the data in the thermodynamic database (pitzer.dat) provided by PHREEQC (Version 3.7.3) released by the U.S. Geological Survey as free software, the values of K10, K20, and KW in eqs. (20), (22), and (24) were calculated.12,13)
2.4 Activity coefficient and H2O activityThe activity coefficients of the ions and the neutral dissolved species and the activity of H2O in the Aq phase were calculated to use the Pitzer model, which is available for concentrated salt solutions. This model considers (1) long-range interactions between ions (using the Debye-Hückel model) and (2) short-range interactions such as two-ion interactions and three-ion interactions as excess Gibbs free energy, GEx = RT ln(γ).
| \begin{align} \frac{G^\text{Ex}}{M_{\text{w}}RT} &= f (I) + \sum_{i}\sum\limits_{j} \lambda_{ij}m_{i}m_{j} \\ &\quad + \sum\limits_{i}\sum\limits_{j}\sum\limits_{k} \mu_{ijk}m_{i}m_{j}m_{k}. \end{align} | (40) |
Here, f(I) is the Debye-Hückel model formula; λij is the second virial coefficient, which is a function of the ionic strength, I; and μijk is the third virial coefficient, which is a function independent of the ionic strength. The Pitzer model can be calculated based on eq. (40). The activity coefficients of γM for cations, γX for anions, and γN for neutral dissolved species (CO2(aq)) used in this study were calculated using the following equations.
| \begin{align} \ln (\gamma_{\text{M}}) & = z_{\text{M}}^{2}F + \sum_{\text{a} = 1}^{N_{\text{a}}} m_{\text{a}} (2B_{\text{Ma}} + ZC_{\text{Ma}}) \\ &\quad + \sum_{\text{c} = 1}^{N_{\text{c}}} m_{\text{c}}\left(2\varPhi_{\text{Mc}} + \sum_{\text{a} = 1}^{N_{\text{a}}} m_{\text{a}}\psi_{\text{Mca}}\right)\\ &\quad + \sum_{\text{a} = 1}^{N_{\text{a}} - 1}\sum_{\text{a}' = \text{a} + 1}^{N_{\text{a}}} m_{\text{a}} m_{\text{a}'}\psi_{\text{aa${'}$M}} + |z_{\text{M}}|\sum_{\text{c} = 1}^{N_{\text{c}}} \sum_{\text{a} = 1}^{N_{\text{a}}} m_{\text{c}}m_{\text{a}}C_{\text{ca}} \end{align} | (41) |
| \begin{align} \ln (\gamma_{\text{X}})& = z_{\text{X}}^{2}F + \sum_{\text{c} = 1}^{N_{\text{c}}} m_{\text{c}}(2B_{\text{cX}} + ZC_{\text{cX}}) \\ &\quad + \sum_{\text{a} = 1}^{N_{\text{a}}} m_{\text{a}}\left(2\varPhi_{\text{Xa}} + \sum_{\text{c} = 1}^{N_{\text{c}}} m_{\text{c}}\psi_{\text{Xac}}\right)\\ &\quad + \sum_{\text{c} = 1}^{N_{\text{c}} - 1} \sum_{\text{c}' = \text{c} + 1}^{N_{\text{c}}} m_{\text{c}}m_{\text{c}'}\psi_{\text{cc${'}$X}} + |z_{\text{X}}| \sum_{\text{c} = 1}^{N_{\text{c}}}\sum_{\text{a} = 1}^{N_{\text{a}}} m_{\text{c}}m_{\text{a}}C_{\text{ca}} \end{align} | (42) |
| \begin{equation} \ln(\gamma_{\text{N}}) = \sum_{\text{c} = 1}^{N_{\text{c}}} 2m_{\text{c}}\lambda_{\text{Nc}} + \sum_{\text{a} = 1}^{N_{\text{a}}} 2m_{\text{a}}\lambda_{\text{Na}} + \sum_{\text{c} = 1}^{N_{\text{c}}} \sum_{\text{a} = 1}^{N_{\text{a}}}m_{\text{c}}m_{\text{a}}\xi_{\text{Nca}} \end{equation} | (43) |
Here, F is the Debye-Hückel term, m is the molality of each ion, zX and zM are the number of charges of the ion, and Z is Σmi|zi|. B, C, Φ, ψ, λ, and ξ are Pitzer parameters or coefficients calculated from Pitzer parameters. The coefficients consist of the temperature function given by the data in the thermodynamic database (pitzer.dat). λ and ξ for CO2–Na+ and CO2–Cl− were obtained from the temperature and pressure functions and data from Duan et al.14) The activity of H2O, aH2O, was calculated here using the formula for the osmotic coefficient ϕosmotic.
| \begin{equation} \ln(a_{\text{H2O}}) = -\phi_{\text{osmotic}}\sum_{n = i} m_{i}/M_{\text{w}} \end{equation} | (44) |
| \begin{align} \phi_{\text{osmotic}} - 1 & = \frac{2}{\displaystyle\sum\nolimits_{n = i}m_{i}}\Biggl[\frac{-A_{\varphi}I^{\frac{3}{2}}}{(1 + bI^{\frac{1}{2}})} \\ &\quad + \sum_{\text{c}}\sum_{\text{a}} m_{\text{c}}m_{\text{a}}(B_{\text{ca}}^{\varphi} + ZC_{\text{ca}})\\ &\quad + \sum_{\text{c}<} \sum_{\text{c}'} m_{\text{c}}m_{\text{c}'}\left(\varPhi_{\text{cc}'}^{\varphi} + \sum_{\text{a}}m_{\text{a}}\psi_{\text{cc${'}$a}}\right)\\ &\quad + \sum_{\text{a}<} \sum_{\text{a}'} m_{\text{a}}m_{\text{a}'}\left(\varPhi_{\text{aa}'}^{\varphi} + \sum_{\text{c}}m_{\text{c}}\psi_{\text{caa}'}\right) \Biggr] \end{align} | (45) |
The detailed-calculation method for these equations is explained in many documents. The method described in Alhseinat et al.15) was used in this study.
The average activity coefficient of a combination of cations and anions was experimentally obtained, and the activity coefficients of the cations and anions were obtained from this coefficient. In an ideal solution, the thermodynamic quantity of each ion is determined based on the reference value of the H+ ion; however, there is no reference in real solutions. The K+ and Cl− ions in a KCl solution have almost the same size and the same mobility, and their activity coefficients are nearly the same. Accordingly, the MacInnes assumption adopted in PHREEQC was applied. This assumption is a correction to match the activity coefficient of the Cl− ion in the solution to be estimated. The following equation was used to correct the activity coefficient for each ion i:16)
| \begin{equation} \gamma_{\text{c}i} = \gamma_{i}\times(\gamma(\mathrm{MA}))^{z_{i}},\quad \gamma(\mathrm{MA}) = \gamma(\text{Cl$^{-}$})/\gamma \pm (\text{KCl}), \end{equation} | (46) |
where γci is the corrected activity coefficient, γ(Cl−) is the activity coefficient of the Cl− ion in the solution to be estimated and γ ± (KCl) is the mean activity coefficient of the KCl solution with equivalent ionic strength.
The values of the temperature and pressure gauges installed at the bottom of the injection wells in the large-scale CCS demonstration test in Tomakomai were 309–321 K and 9.28–10.07 MPaG in the Moebetsu Formation (an aquifer of sandstone, vertical depth of 1188 m), and approximately 360 K and 22.84–37.07 MPaG in the Takinoue Formation (an aquifer of volcanic rocks, vertical depth of 2753 m).17) Table 1 shows an example of a formation water analysis during exploratory drilling of the Moebetsu Formation. It is suggested that the Moebetsu Formation water has characteristics of ancient seawater containing a large amount of fresh water. The analysis was conducted on the as-collected formation water, which was low in salinity and high in pH and contained HCO3− ions. It is generally believed that formation water originates from seawater. Formation water with higher Cl− ion concentrations than seawater is relating to (1) evaporation resulting from temperature increases, (2) separation by specific gravity, (3) adsorption by clay minerals, (4) consolidation action, and/or (5) semi-permeable membrane effects caused by clay minerals etc.; in addition, inter-layer water resulting from the phase transition of clay minerals is cited as causing decreases in the Cl− ion concentration.18) Ions such as metal ions of HCO3− and SO42− change because of reactions with minerals, aging of organic matter and volcanic gases. Ahmadun et al. reviewed the chemical composition of produced water in oil and gas wells (considered to have the same generation factor as formation water), finding that some environments had salinity levels as high as 190,000 mg/L.19) Following the injection of high-pressure CO2, CO2 can temporarily exist in an aquifer in a free state; however, the formation water gradually returns and dissolves the CO2. In a supercritical CO2 environment, CO2 is expected to transform into carbonate rocks over a long period of time. Long-term CO2 corrosion then becomes a problem. Further considering the transport of supercritical CO2, the environment of deep wells, and enhanced oil recovery (EOR), the validation of the U-Cal model was performed to compare measured values in the literature with calculation values in the U-Cal model, of environmental factors in pure water to high-concentration NaCl solutions at temperatures from room temperature to 423.15 K and in pressures up to 50 MPa.
The data compiled by Spycher et al. and the data by Ahmadi et al. were used for the experimental values of the CO2 solubility m(CO2) in the Aq phase.20,21) Table 2 shows the average absolute relative deviation, AARD (%), calculated from the mesdured and calculated values, AARD (%) = 100/Nexp Σ(|Xexp − Xcal|/Xexp), where Nexp, Xexp, and Xcal are the number of data points, experimental value, and calculated value, respectively. AARD is 1.7%–2.89% when the measured values, which deviate greatly, are excluded. These values indicate that the calculation results derived by the U-Cal model are in good agreement with the measured values. Figure 1 shows the pressure and temperature dependence of m(CO2). At 298.15 K, which is below the critical point, m(CO2) is larger at any pressure than at higher temperatures than 298.15 K. At this temperature, CO2 becomes gas, liquid + gas, and then subcritical with increasing pressure. In the gaseous state, m(CO2) increases remarkably as the pressure increases. At approximately 2–3 MPa, m(CO2) is nearly equal to the value calculated by the ideal gas-ideal solution model. In the pressure (total pressure) more than the liquefaction pressure, the state changes to liquid + gas and subcritical CO2, and the rate of increase of m(CO2) decreases. At temperatures above the critical temperature of CO2, m(CO2) increases gradually with increasing temperature. Relevant to storage in an aquifer, the CO2 storage efficiency is very high at approximately 10 MPa at temperatures close to 323 K because m(CO2) is very large under those conditions. At pressures above 30 MPa, m(CO2) does not change much between 323 K and 423 K.
Effect of total pressure and temperature on CO2 solubility in aqueous (Aq) phase in CO2–H2O system.
The molar fraction of H2O, y(H2O), in the Naq phase is larger than that in the ideal gas in the supercritical state because of the mutual dissolution of CO2 and H2O as shown in Fig. 2. At 298.15 K, like m(CO2) in Fig. 1, y(H2O) behaves like an ideal gas under gaseous CO2 state; however, it suddenly increases when the state became liquid + gas and then remains nearly unchanged even if the pressure is increased. In the Naq phase (CO2-rich phase), the water content y(H2O)/g L−1, grams of water in 1 L of the Naq phase, is easy to visualize when considering corrosion. The water content is plotted on the vertical axis in Fig. 3. At 298.15 K, the pressure increases and the water content increases abruptly in the region where liquid CO2 forms, with higher temperatures leading to greater water contents. At any temperature with the total pressure not less than 15 MPa, this change becomes small; however, 1–10 g/L of water still exists. Therefore, corrosion resulting from relatively high amounts of water should be evaluated in the transport of supercritical CO2.
Effect of total pressure and temperature on water mole fraction in non-aqueous (Naq) phase in CO2–H2O system.
Effect of total pressure and temperature on water solubility in non-aqueous (Naq) phase in CO2–H2O system.
The pressure and temperature dependence of the pH in the Aq phase is shown in Fig. 4 with the experimental values of Peng et al.22) Above 6 MPa, the pH value is nearly constant regardless of the temperature. At 308.3 K, which is a temperature close to the critical state in the Naq phase, the calculated pH is approximately 0.1 higher than the measured value for pressures above 6 MPa. Peng et al. used a glass electrode with effective operating temperatures of 273–353 K and zirconia electrode with effective operating temperatures of 363–578 K as pH electrode, and a Ag/AgCl electrode as a reference electrode. The uncertainty in the measured pH value estimated with 95% confidence was 0.06 pH units. Because this value corresponds to the absolute relative deviation of 1.62% at pH 3.7, it is nearly equal to an AARD of 1.7% shown in Table 2. It would be difficult to measure and estimate environmental factors near the critical point. As the pressure increases from a low pressure of approximately 0.5 MPa to approximately 6 MPa or higher, the pH drops by approximately 0.55 at 308.3 K and by approximately 0.85 at 423.15 K.
Effect of total pressure and temperature on pH in CO2–H2O system.
Figure 5 shows the effect of the HCO3− ion concentration and pressure on the pH at 353.15 K for comparison with the experimental values reported by Li et al.23) Although the pH depends greatly on the HCO3− ion concentration, it does not decrease significantly above 5 MPa. For example, at 10 MPa, the predicted pH values in the U-Cal model are 3.19 at 0 mol/kg, 3.49 at 0.001 mol/kg, 4.35 at 0.01 mol/kg, and 5.29 at 0.1 mol/kg. AARD in Table 2 is 2.77%, and the predicted pH value agrees well with the experimental value. However, when the HCO3− ion concentration is as high as 0.1 mol/kg, AARD increases to approximately 6% at pressures above 10 MPa.
Effect of total pressure and HCO3− ion concentration on pH at 353.15 K in CO2–H2O system.
Figure 6 shows the values of m(CO2) measured by Zhao et al.24) at a total pressure of 15 MPa and the values calculated to use the U-Cal model. When the NaCl concentration increases from pure water to 6 mol/kg, m(CO2) becomes approximately 1/3 at 298.15 K and approximately 1/2.5 at 423.15 K. However, the activity of CO2(aq), a(CO2), is hardly affected by the NaCl concentration. For the corrosion reaction, the lowering of the CO2 solubility in a high-concentration NaCl solution means that the ability to supply H+ ions by eq. (19) decreases. Figure 7 shows the NaCl concentration and temperature dependence of the pH. The experimental values of the pH are taken from Li et al.23) Li et al. measured the pH using the same electrode configuration as Peng et al.,22) and the uncertainty in the pH at 95% reliability was 0.2 pH units, which is larger than that of pure water. The difference between the measured and calculated pH values is larger at 308.3 K than at other temperatures. This is thought to be due to difficulties in measuring the pH near the critical point and in predicting the activity coefficient. The pH decreases by approximately 0.55 when the NaCl concentration increases from pure water to 5 mol/kg.
Effect of NaCl molarity on CO2 solubility in aqueous (Aq) phase under 15 MPa total pressure in CO2–H2O–NaCl system.
Effect of NaCl molarity on pH under 15.38 MPa total pressure in CO2–H2O–NaCl system (*Data at 0 m NaCl is from Peng et al.22)).
The main features of carbon steel corrosion in CO2 environments are the increasing hydrogen ion supply power in the cathodic reaction and the dissolution of Fe and the formation of FeCO3 in the anodic reaction. Figure 8 shows the potential-pH diagram at 333.15 K in the Fe–CO2–H2O system in 0.1 MPa and 15 MPa (fugacity, f = 7.9 MPa) CO2 environments with a Fe2+ ion concentration of 10−4 mol/L. It is suggested that Fe dissolves as Fe2+ ions at the initial stage of corrosion in pH regions of CCS environments and that FeCO3 is produced as the corrosion progresses and the pH increases.
Potential-pH diagram at 333.15 K in Fe–CO2–H2O system (CO2 fugacity: 7.9 MPa (15 MPa total pressure) and 0.1 MPa, Fe2+ ion: 10−4 mol/L, other ions: 10−6 mol/L, H2: 0.1 MPa).
Waad et al. electrochemically studied CO2 corrosion at CO2 partial pressures up to 0.1 MPa and temperatures up to 353.15 K, and proposed the following corrosion rate equation;25)
| \begin{equation} \log(\textit{CR}) = -A \times \text{pH} + B, \end{equation} | (47) |
where CR is the corrosion rate, A is 1.3 and B is a constant. A large value of A is characteristic of CO2 corrosion. This A can be explained by considering the pH-dependent dissolution reaction of iron and the cathodic reaction by the direct reduction reaction of H2CO3(aq) in the following equation;
| \begin{equation} \text{H$_{2}$CO$_{3}$(aq)} + \text{e$^{-}$} \to \text{H} + \text{HCO$_{3}{}^{-}$}. \end{equation} | (48) |
H2CO3(aq) is supplied from the diffused H2CO3(aq), H+ ions and HCO3− ions. The relationship between the corrosion rate calculated from the weight loss of carbon steel examined by Choi et al.26) and the pH calculated by the U-Cal model was investigated, in 48-h tests in 25% NaCl + CO2 solutions at 338.15 K with total pressures of 4, 8, and 12 MPa. In these environments, the corrosion rates obtained electrochemically by Choi et al.26) were nearly constant from 10 h to approximately 40 h. Then a small amount of FeCO3 and the residual substance in the steel, Fe3C, were only present locally on the steel after tested. It is thought that these environments primarily experienced the dissolution reaction of iron. Based on the corrosion rate at 4 MPa, the corrosion rates, that is, the dissolution rates, at 8 and 12 MPa using eq. (47) were calculated such that
| \begin{equation} \frac{\textit{CR}(x\,\text{MPa})}{\textit{CR}(\text{4$\,$MPa})} = 10^{-A(\text{pH(${x}{}\,$MPa)} - \text{pH(4$\,$MPa)})}. \end{equation} | (49) |
Figure 9 shows the measured and calculated corrosion rates, which show good agreement. This suggests that the iron dissolution reaction is a function of the pH, even in high-pressure CO2 environments. Because H2CO3(aq) is a weak acid, the pressure dependence of the buffering action was also investigated. In the same environments as shown in Fig. 9, the changes in the pH with 100 ppm dissolved Fe2+ ions (ΔpH = pH at 100 ppm Fe2+ − pH at 0 ppm Fe2+) were calculated to use the U-Cal model. As shown in Fig. 10, a higher pressure corresponds to a higher m(CO2) concentration, smaller ΔpH value, and larger buffering capacity. The pit penetration rate measured by Choi et al. had a large value in a CO2 environment of 12 MPa as shown in Fig. 9, because localized corrosion is also a problem in CO2 corrosion.
Effect of total pressure on corrosion rate and pit penetration rate in CO2–H2O–NaCl system26) (338.15 K, 25 mass% NaCl solution, carbon steel, 48 h).
Effect of total pressure on ΔpH and CO2 solubility in CO2–H2O–NaCl system (338.15 K, 25 mass% NaCl, ΔpH = pH at 100 ppm Fe2+ − pH at 0 ppm Fe2+).
The Fe2+ ion concentration, meq(Fe2+), under the equilibrium state of FeCO3 formation was calculated using the law of conservation of mass and the condition of electrical neutrality and adding the following reactions to eqs. (19), (21) and (23).
| \begin{equation} \text{Fe} + \text{2H$^{+}$} = \text{Fe$^{2+}$} + \text{H$_{2}$} \end{equation} | (50) |
| \begin{equation} \text{Fe$^{2+}$} + \text{CO$_{3}{}^{2-}$} = \text{FeCO$_{3}$},\quad m(\text{Fe$^{2+}$})m(\text{CO$_{3}{}^{2-}$}) = K_{\text{sp}} \end{equation} | (51) |
| \begin{align} \log(K_{\text{sp}}) & = -59.3498 - 0.041377 \times T - 2.1963/T \\ &\quad+ 24.5724\times \log(T)\\ &\quad + 2.518 \times I^{0.5} - 0.657\times I \end{align} | (52) |
The formula obtained by Sun et al. was used for Ksp.27) The term considering the ionic strength I is a relational expression obtained from the aqueous solution with I = 0.1–5.5 mol/kg at 298.15 K according to Silva et al.28) The condition for electrical neutrality is as follows:
| \begin{align} &m(\text{H$^{+}$}) + 2m(\text{Fe$^{2+}$}) \\ &\quad = m(\text{OH$^{-}$}) + m(\text{HCO$_{3}{}^{-}$}) + 2m(\text{CO$_{3}{}^{2-}$}). \end{align} | (53) |
Using the equilibrium constants of eqs. (29)–(32) and (52), this eq. (53) is converted to the following eq. (54) for m(H+) only; m(H+) is then obtained by the iterative method.
| \begin{align} \mathit{func}(m(\text{H$^{+}$})) & = 2K_{\text{sp}}/(f(\text{CO$_{2}$})K_{\text{B}})m(\text{H$^{+}$})^{4}\\ & \quad + m(\text{H$^{+}$})^{3} - (K_{\text{W}} + f(\text{CO$_{2}$}) K_{\text{A}})m(\text{H$^{+}$})\\ &\quad - 2f(\text{CO$_{2}$})K_{\text{B}}\\ &= 0 \end{align} | (54) |
Then, meq(Fe2+) is calculated using eqs. (30) and (51).
| \begin{equation} m_{\text{eq}}(\text{Fe$^{2+}$}) = K_{\text{sp}} m(\text{H$^{+}$})^{2}/(f(\text{CO$_{2}$})K_{\text{B}}) \end{equation} | (55) |
Figure 11 shows the temperature and pressure dependence of meq(Fe2+) in pure water. meq(Fe2+) decreases with increasing temperature. That is, if the total pressure is the same, it is easier to produce FeCO3 at a higher temperature. As for the pressure dependence, meq(Fe2+) becomes nearly constant above 10 MPa. When gaseous CO2 exists in the Naq phase, the temperature dependence of meq(Fe2+) is nearly the same as that calculated from the FeCO3 formation model (model 2) considering the pH change resulting from corrosion obtained by Ueda.29) Figure 12 shows the temperature and NaCl concentration dependence of meq(Fe2+) at a total pressure of 15 MPa. meq(Fe2+) increases with increasing NaCl concentration up to approximately 1 mol/kg, but decreases at higher NaCl concentrations. This decreasing tendency becomes smaller as the temperature increases.
Effect of total pressure and temperature on Fe2+ concentration under equilibrium condition of FeCO3 formation in Fe–CO2–H2O system.
Effect of NaCl morality and temperature on Fe2+ concentration under equilibrium condition of FeCO3 formation in Fe–CO2–H2O–NaCl system (Total pressure: 15 MPa).
According to the experimental results of Ikeda et al., the Fe2+ ion concentrations m(Fe2+) measured after a 96-h corrosion experiment of pure iron in environments with 3.0 MPa CO2 and 5% NaCl from ambient temperature to 423.15 K were larger than the meq(Fe2+) values in the equilibrium state of FeCO3 formation.30) The results of time change in measured m(Fe2+) of AISI 4140 (carbon steel) in environments with 15.6 MPa CO2 and 1650 ppm NaCl at 355 K for up to approximately 20 h obtained by Dunlop et al. were also much larger than the meq(Fe2+) values.31) The relationship between the time dependence of the Fe2+ ion concentration, m(Fe2+), and the corrosion behavior of carbon steel was investigated in detail by Dugstad in a pure water + CO2 environment with a total pressure of 0.1 MPa.32) Because water can accumulate up to approximately 4 cm in a horizontal large-diameter line pipe, the experiments were conducted under a very small specific volume/area ratio (that is, the ratio of the test solution to the total surface area of the test piece) of 4 mL/cm2. The obtained-results are summarized in Fig. 13.
1st stage: Fe dissolution; 2nd stage: Fe dissolution and precipitation of FeCO3 from the supersaturated state; and 3rd stage: stages 1 and 2 become steady states.
Effect of time and temperature on Fe2+ concentration in CO2–H2O system32) (Total pressure: 0.1 MPa, pure water, Material: carbon steel).
The measured concentration of Fe2+ ion, when a FeCO3 was deposited on AISI 316 specimen and then the specimen immersed in a CO2 environment with total pressure of 0.1 MPa, was nearly the same as the value of meq(Fe2+) obtained by the equilibrium calculation.32) That is, once FeCO3 forms, it dissolves only up to meq(Fe2+), which is the calculated equilibrium value.
The FeCO3 crystal growth rate (PR) formula, which can be applied to a wide range of supersaturation (SS), was proposed by Sun et al.33) We considered applying the following equation to the time change of Fe ion concentration under FeCO3 formation:
| \begin{align} \frac{-dm(\text{Fe$^{2+}$})}{dt} &= \textit{PR} = e^{(A - \frac{B}{RT})}\frac{S}{V}K_{\text{sp}} (\textit{SS} - 1) \\ &= \textit{PR}_{0}(\textit{SS} - 1), \end{align} | (56) |
| \begin{equation*}\textit{SS} = \frac{m(\text{Fe$^{2+}$})m(\text{CO$_{3}^{2-}$})}{K_{\text{sp}}}, \end{equation*} |
where A = 28.2 kJ/mol, B = 64.850 kJ/mol, S is the total surface area of the specimen (m2), and V is the volume of the test solution (m3). Dugstad assumed that (SS-1) was equivalent to the relative supersaturation (RS = (m(Fe2+) − meq(Fe2+))/meq(Fe2+)).34) Adopting this assumption, the FeCO3 crystal growth rate formula is as follows.
| \begin{equation} \frac{-dm(\text{Fe$^{2+}$})}{dt} = \textit{PR} _{0}\left(\frac{m(\text{Fe$^{2+}$})}{m_{\text{eq}}(\text{Fe$^{2+}$})}-1 \right). \end{equation} | (57) |
Integrating eq. (57), m(Fe2+) can be obtained such that
| \begin{align} m(\text{Fe$^{2+}$})& = m_{\text{eq}}(\text{Fe$^{2+}$}) + (m_{0}(\text{Fe$^{2+}$}) \\ &\quad - m_{\text{eq}}(\text{Fe$^{2+}$})) \exp \left(-\frac{\textit{PR} _{0}}{m_{\text{eq}}(\text{Fe$^{2+}$})}t \right), \end{align} | (58) |
where m0(Fe2+) is the Fe2+ ion concentration at 0 h. Figure 14 shows the calculation results for m(Fe2+) in the same CO2 gas + pure water environment as that of Dugstad’s experiment shown in Fig. 13. The concentration at which the m(Fe2+) in Fig. 13 begins to decrease is defined as the Fe2+ concentration m0(Fe2+) at 0 h. The relative supersaturations (RS) at 0 h were estimated to be 13 at 293.15 K, 8 at 333.15 K, and 6 at 353.15 K. In addition, the RS value at 423.15 K with a CO2 fugacity of 0.1 MPa (a total pressure 0.571 Ma) was assumed to be 3 and the calculation result for m(Fe2+) is shown for reference.
Effect of time and temperature on Fe2+ concentration in Fe–CO2–H2O system (Total pressure: 0.1 MPa, pure water, Specific volume: 4 mL/cm2).
Corrosion of carbon steel in a high-pressure CO2 environment is regarded as pH-dependent corrosion, as discussed for the dissolution of iron, in an environment where the corrosion product FeCO3 does not occur. However, in an environment where FeCO3 produces, it is necessary to discuss the state in which Fe2+ ions dissolve in supersaturation, that is, the metastable state, which is a difficult problem. Using the experimental result of Choi et al.26) in a high-pressure CO2 environment with 25% NaCl at 363.15 K, we considered corrosion in an environment where FeCO3 produces. Figure 15 shows the dependence of total pressure and time on the electrochemically obtained corrosion rate.
Effect of time and total pressure on corrosion rate in CO2–H2O–NaCl system26) (363.15 K, 25 mass% NaCl, Material: carbon steel).
Effect of time and total pressure on Fe2+ concentration calculated from eq. (58) and meq(Fe2+) in CO2–H2O–NaCl system (25 mass% NaCl, Specific volume: 4 mL/cm2).
Choi et al. also observed a maximum pit penetration rate of 19 mm/y as a result of localized corrosion.26)
Temperature dependence of the corrosion rate
Figure 17 shows the corrosion rate of carbon steel in a supercritical CO2 + pure water environment at 9.5 MPa total pressure obtained by Zhang et al.37) for comparison with the corrosion rate of pure iron in a 5% NaCl solution environment with 3.0 MPa CO2 (at room temperature) measured by Ueda.29) The temperature Tmax showing the maximum corrosion rate was nearly the same, even if the state of the CO2 in the Naq phase differed from a supercritical fluid to a gas. It is thought that at temperatures below Tmax, iron dissolution dominant corrosion occurred, while at temperature above Tmax, FeCO3 formation-dominant corrosion occurred. According to the U-Cal model, the pH in Zhang et al.’s experimental environment was approximately 0.1 lower than that in Ueda’s experimental environment.
Effect of HCO3− ion concentration
Kahyarian et al. investigated the effects of pH on the corrosion rate of carbon steel (API X65) by the addition of NaHCO3 at 283.15 K and 313.15 K in 0.1 M NaCl solutions with 0.5 MPa CO2 from the cathodic and anodic polarization curves.38) The anodic polarization curve shows the dissolution behavior of iron. At 283.15 K, the corrosion rate was nearly the same (approximately 0.6 mm/y) at pH 4 and 5 and decreased slightly to approximately 0.5 mm/y at pH 6. At 313.15 K, the corrosion rate was approximately 15 mm/y at pH 4, approximately 6 mm/y at pH 5, and approximately 4 mm/y at pH 6. This result is thought to be related to the hydrogen ion supplying ability because of the buffering action of carbonic acid, which is a weak acid. The cathodic polarization curve showed a plateau corresponding to the limiting current of the hydrogen ions (the current density is constant regardless of the potential), and the pH dependence was small. Ikeda et al. obtained the current density at −0.65 V versus SCE (saturated calomel electrode) corresponding to this plateau from the cathodic polarization curve of platinum in 0.1 MPa CO2 at 298.15 K and 5% NaCl environment. It was approximately 100 µA/cm2 at pH 4, decreased by approximately 20% at pH 5, and showed nearly the same value at pH 6 as at pH 5.39) These results indicate that the corrosion rate is correlated with the limiting current density of the hydrogen ions in a CO2 environment in which iron is dissolved, such that, even if the HCO3− ion concentration increases and the pH rises significantly, the corrosion rate will not change significantly. In addition, in an environment where FeCO3 produces, corrosion is thought to be suppressed because FeCO3 is more likely to form on the steel surface when the concentration of HCO3− ions increases.
4.5 Corrosion of carbon steel in the Naq phaseIn the supercritical CO2 state, the water fraction of the Naq phase is much larger than that expected from the water vapor pressure in the ideal state, as shown in Fig. 2, because of the mutual dissolution of CO2 and H2O. 96 h corrosion tests of carbon steel simulating the environment of the Naq phase were conducted by Zhang et al. to use a two-liter autoclave under the test conditions of a total pressure of 6.0–12.6 MPa, a temperature of 323.15–403.15 K, 1000 ppm NaCl:40) an accelerated corrosion test in the Naq phase with 600 g CO2 + 100 g H2O.
A rotating cylinder, in which the test specimens were set, was placed in the Naq phase, and a propeller was attached to the bottom of the rotating body to stir the Aq phase. The specimen was in a mist water solution + CO2 (gas or supercritical fluid) environment.
The results are shown in Fig. 18, and the corrosion rates were very small, between 14 and 43 µm/y. In addition, localized corrosions of 13–21 µm were observed. The corrosion rate was maximum at 353.15 K. This temperature dependence on the corrosion rate in Naq phase is the same as that in the Aq phase, suggesting a relationship with the FeCO3 formation behavior. Thodla et al. electrochemically measured the corrosion rate of carbon steel in a supercritical CO2 environment (304.15 K and 7.93 MPa) containing 100 and 1000 ppm H2O and found large corrosion rate of approximately 2 mm/y in 100 ppm H2O and approximately 1 mm/y in 1000 ppm H2O;41) this is thought to correspond to the high corrosion rate in the initial stage of corrosion, as discussed in the case of corrosion in the Aq phase.
Effect of temperature on corrosion rate in CO2–H2O–NaCl system40) (600 g CO2 + 100 g H2O, 1000 ppm NaCl, 96 h, carbon steel).
A thermodynamic model (U-Cal model) was proposed to predict the fugacity of CO2 and the water content in the CO2-rich phase (gas or supercritical fluid), as well as the CO2 solubility and pH in aqueous solutions in CCS environments, that is, a CO2 environment in a supercritical state. The CO2 corrosion behavior of carbon steel was discussed to use the U-Cal model. The following conclusions were obtained.
We would like to express our gratitude to the Tomakomai CCS Association for allowing us to use their valuable CCS data. We would also like to express our sincere gratitude to the individuals at the technical research center of the INPEX corporation who provided useful advice on geology and the behavior of CO2 in underground environments.
