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Borate-Catalyzed Carbon Dioxide Hydration via the Carbonic Anhydrase Mechanism Dongfang Guo,†,‡ Hendy Thee,‡ Gabriel da Silva,*,‡ Jian Chen,† Weiyang Fei,† Sandra Kentish,‡ and Geoffrey W. Stevens‡ † ‡
State Key Laboratory of Chemical Engineering, Department of Chemical Engineering, Tsinghua University, Beijing 100084, China Department of Chemical and Biomolecular Engineering, The University of Melbourne, Victoria 3010, Australia
bS Supporting Information ABSTRACT: The hydration of CO2 plays a critical role in carbon capture and geoengineering technologies currently under development to mitigate anthropogenic global warming and in environmental processes such as ocean acidification. Here we reveal that borate catalyzes the conversion of CO2 to HCO3 via the same fundamental mechanism as the enzyme carbonic anhydrase, which is responsible for CO2 hydration in the human body. In this mechanism the tetrahydroxyborate ion, B(OH)4, is the active form of boron that undergoes direct reaction with CO2. In addition to being able to accelerate CO2 hydration in alkaline solvents used for carbon capture, we hypothesize that this mechanism controls CO2 uptake by certain saline bodies of water, such as Mono Lake (California), where previously inexplicable influx rates of inorganic carbon have created unique chemistry. The new understanding of CO2 hydration provided here should lead to improved models for the carbon cycle in highly saline bodies of water and to advances in carbon capture and geoengineering technology.
1. INTRODUCTION The hydration of aqueous carbon dioxide (CO2) to the bicarbonate ion (HCO3) plays a crucial role in emerging technologies that are being actively pursued to mitigate rising atmospheric carbon dioxide levels, such as the absorption of CO2(g) into aqueous solvents for carbon capture and storage (CCS), carbon dioxide air capture, and other geoengineering schemes.14 Furthermore, CO2 hydration is a fundamental reaction critical to environmental processes such as ocean acidification and CO2 mineralization in the global carbon cycle.57 The CO2/HCO3 hydration/dehydration cycle also plays an important role in pH buffering in the human body, where these relatively slow chemical reactions are facilitated by the enzyme carbonic anhydrase.8 Because of the importance of CO2 hydration, this process has been thoroughly investigated by many researchers.811 It is well established that in aqueous media this process is facilitated by both H2O and OH. In acidic to neutral solutions, the following reaction takes place (in this work all species are aqueous unless otherwise stated): CO2 þ H2 O f H2 CO3
ð1Þ
The subsequent deprotonation of carbonic acid (H2CO3) to the bicarbonate ion (HCO3) þ Hþ is rapid, and as such the overall reaction is often written as CO2 þ H2O f HCO3 þ Hþ. Further deprotonation of HCO3 produces the carbonate ion r 2011 American Chemical Society
(CO32-) þ Hþ. As pH increases, the reaction of CO2 with the hydroxide ion (OH) begins to dominate the aqueous hydration process: CO2 þ OH f HCO3
ð2Þ
It is well-known that CO2 hydration can be catalyzed by inorganic species such as arsenate and borate.1214 The discovery of new compounds that accelerate CO2 hydration could lead to novel solvents for carbon capture and improved geoengineering technologies. Given the diverse range of chemicals dissolved in saline bodies of water, previously unconsidered catalysts may be responsible for accelerated rates of atmospheric CO2 absorption.15 Several studies on the reactive absorption of CO2 into alkaline solvents have used boric acid to accelerate the rate of reaction.13,14 However, the mechanism involved is not understood and it is uncertain what the active form of the boron compound is and whether or not it chemically binds CO2 (as is the case for commonly used amine solvents16) or actually catalyzes CO2 hydration. Given that boron will be present in both the boric acid and borate forms at low boron concentrations across the pH range relevant to environmental processes and carbon capture systems,17,18 the following reactions need Received: February 21, 2011 Accepted: April 15, 2011 Revised: April 14, 2011 Published: May 02, 2011 4802
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to be considered:
CO2 þ BðOHÞ4 f products
ð3Þ
CO2 þ BðOHÞ3 f products
ð4Þ
The purpose of the present work is to investigate the reaction kinetics and mechanism of boron-catalyzed CO2 hydration using stopped-flow UV/visible spectrophotometry. Specifically, we consider the possibility of CO2 reacting with B(OH)4 (eq 3) and B(OH)3 (eq 4). Subsequent to this we explore the potential impact of this chemistry in carbon capture processes and in environmental systems where elevated levels of boron are encountered.
2. MATERIALS AND METHODS All reagents employed here were of analytical reagent grade. Two indicators and one buffer were purchased and used without further purification: 4-nitrophenol (g99.5%, Sigma), thymol blue (Sigma), and imidazole (g99.0%, Sigma). Hydrochloric acid standard solution (1 M, Merck) and sodium hydroxide (g97.0%, Chem-Supply) solutions were used to adjust pH values. The potential catalyst boric acid (g99.5%) was purchased from Chem-Supply. CO2/N2 gas mixtures (10.6% and 14.0% CO2 in N2 obtained from BOC Gases Australia Limited) were used for the preparation of aqueous CO2 solutions. The stopped-flow pH indicator technique was adopted to study both uncatalyzed and catalyzed CO2 hydration reactions. Reaction was initiated by mixing a solution containing CO2(aq) with a reagent solution containing pH indicator and any added buffer and/or catalyst. CO2 solutions were freshly prepared for each batch of experiments by bubbling the CO2/N2 gas mixture through a gas absorption bottle containing deionized, distilled water for at least one hour before the experiment began, with the flow of gas continued as long as the solution was in use. A gastight syringe with a needle was used to withdraw and contain the CO2 solution, which was immediately connected to the sample inlet of the stopped-flow apparatus to avoid the solution contacting air. The CO2(aq) concentration was determined from published solubility data, and was typically in the range of 1.6 to 2.1 mM in our experiments. Reagent solutions were prepared using deionized, distilled water and pH adjusted by adding hydrochloric acid or sodium hydroxide. The reaction pH was determined by mixing the CO2 and reagent solutions outside of the stopped-flow cell, using a Metrohm Titrando 809 autotitrator (Switzerland). To study the CO2 þ OH reaction, the reagent solutions contained thymol blue (1.18 105 M) indicator, with pH ranging from 8.4 to 9.7. The CO2 þ H2O reaction was studied using imidazole buffer solutions (820 mM, pH = 6.8) with 4-nitrophenol (2.35 105 M) as the indicator. Boroncatalyzed CO2 hydration experiments were performed in the B(OH)3/B(OH)4 buffer system at different pH values (8.89.6), with thymol blue indicator (1.18 105 M) and constant total boron concentration (6.25 mM added boric acid). All quoted concentrations refer to the reaction media after mixing. Kinetic measurements reported here were performed on an Applied Photophysics SX.18MV-R Stopped-Flow Reaction Analyzer equipped with a 10-mm optical path length observation chamber (Applied Photophysics Ltd., United Kingdom). The stopped-flow cell and reagent reservoirs were temperatureregulated to within (0.1 °C using a Grant water bath. Prior to kinetic experiments, spectrophotometric properties for all reagents were measured between 200 and 800 nm in the stopped-flow cell.
Extinction coefficients (ε) for the base and acid forms of the pH indicator thymol blue at 598 nm are 3.57 104 M1 cm1 and 230 M1 cm1, respectively, whereas those for the indicator 4-nitrophenol at 400 nm are measured as 1.81 104 M1 cm1 and 100 M1 cm1. A typical kinetic experiment was initiated by mixing the aqueous reagent solutions in a 1:1 ratio, and monitoring the indicator absorbance versus time using the SX.18MV-R software. Experiments were performed in which both pH and temperature were varied. Seven repeat runs were performed at each set of conditions, and values reported here represent averages. Reaction was typically complete within several seconds, and initial reaction rates (corresponding to around 10% conversion) were obtained from the recorded absorbance traces by a signal exponential regression based on the Marquardt algorithm. Conditions were chosen such that [CO2(aq)] was always in excess over [OH], resulting in pseudo-first-order kinetics. The technique used to extract pseudo-first-order rate constants (kobs, s1) from the initial reaction rates is described in the Supporting Information. Standard deviations from the repeat runs were used to calculate uncertainty intervals for the kobs values, which are twice the standard error. Uncertainties are propagated through the rate constant calculations in the usual manner to provide error estimates for the final intrinsic rate constants.
3. RESULTS The kinetics of uncatalyzed CO2 hydration (i.e., reaction with H2O and OH) was first investigated, as measurements of k1 and k2 are required under our experimental conditions so that we can accurately determine k3 and k4. Additionally, these experiments validate the experimental approach adopted here, and also provide new rate constant measurements for these important reactions. These experiments are described in further detail in the Supporting Information. For the CO2 þ H2O reaction we measure k1 = 0.0298 ( 0.0007 s1 at 25 °C, with rate constants between 25 and 40 °C described by the Arrhenius expression k1 [s1] = 4.07 106 exp(5584/T [K]). For CO2 þ OH we find k2 = 6830 ( 200 M1 s1 with Arrhenius expression k2 [M1 s1] = 9.88 1013 exp(6956/T [K]) from 20 to 40 °C. Our value for the slow CO2 þ H2O reaction rate constant at 25 °C is similar to but somewhat smaller than the value of 0.037 ( 0.002 s1 reported by Khalifah.8 The measured rate constant for the CO2 þ OH reaction at standard temperature compares well with previous measurements of 8500 M1 s1,10 6000 M1 s1,11 and 7900 M1 s1,12 which are both larger and smaller than our value. For the boron-catalyzed CO2 hydration reaction, under our experimental conditions (6.25 mM boron) the total boric acid concentration is sufficiently low that B(OH)3 is only in equilibrium with B(OH)4, and the complicating presence of polyborate ions could be ignored. Total observed rate constants, kobs, are plotted in Figure 2A as a function of [OH]. When the component due to reaction with OH and H2O is removed (k0 obs = kobs k2[OH] k1) we observe nonzero rate constants, which provides evidence for the catalysis of CO2 hydration by the added boron. A plot of k0 obs versus [B(OH)4] (Figure 2B) reveals that the catalyzed component of the rate constant is first-order in tetrahydroxyborate ion concentration, suggesting that this is the active species. Furthermore, when the linear regressions are extrapolated to [B(OH)4] = 0 (i.e., when [B]total ≈ [B(OH)3]) the intercepts yield small negative values 4803
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Environmental Science & Technology (ca. 0.05 s1), demonstrating that B(OH)3 catalysis is not significant. The slope of the linear regression of k0 obs versus [B(OH)4] provides the rate constant k3 for the CO2 þ B(OH)4 reaction. At 25 °C k3 is measured to be 35.3 ( 2.0 M1 s1; from 25 to 40 °C k3 [M1 s1] = 2.20 1013exp (8106/T [K]), where the activation energy is 67.4 ( 2.7 kJ mol1 (from Figure 3) To confirm the active form of boron catalysis, a series of experiments were performed under the same conditions used to measure the CO2 þ H2O rate constant, but now in the presence of boric acid (4.0 mM). These experiments were designed to detect if B(OH)3 can catalyze CO2 hydration (at around pH 6.8 B(OH)3 is the dominant boron compound; cf. Figure 1). No evidence is found for catalysis of CO2 hydration by the neutral boric acid molecule (see Figure S1, Supporting Information).
4. DISCUSSION The ability for boron compounds to accelerate the reactive absorption of CO2 into aqueous media is not unknown, although the mechanism for this process, and the active form of boron,
Figure 1. Distribution (%) of B(OH)4 (versus B(OH)3) in borate solutions in the pH range 513 at 25 °C (solid line) and 60 °C (dashed line).
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remains elusive. Results presented here clearly indicate that B(OH)4 catalyzes the conversion of CO2 to HCO3. This process is first order in [B(OH)4], suggesting that this ion undergoes a direct reaction with CO2. The pre-exponential factor for this reaction, (2.20 ( 0.08) 1013 M1 s1, is relatively large for an elementary solution-phase reaction but is of magnitude similar to that for the CO2 þ OH reaction, (9.88 ( 0.37) 1013 M1 s1. The activation energy of 67.4 ( 2.7 kJ mol1 is somewhat larger than that for both the CO2 þ OH and CO2 þ H2O reactions (measured here as 57.8 ( 2.8 kJ mol1 and 46.4 ( 0.6 kJ mol1, respectively). We propose the mechanism depicted in Scheme 1 for the borate-catalyzed hydration of CO2. Here, the boric acidwater complex deprotonates to form B(OH)4 þ Hþ, in a rapid (equilibrium) reaction.18 The borate ion then reacts with CO2 in what is the rate-determining step, producing a B(OH)3 3 HCO3 intermediate. The HCO3 moiety in this species is replaced by a water molecule, producing HCO3 and regenerating the initial boric acid molecule. This final step is expected to be fast, and may indeed proceed simultaneously with the B(OH)4 þ CO2 reaction in a concerted mechanism, as opposed to the stepwise process described here in which B(OH)3 3 HCO3 exists as a
Figure 3. Arrhenius plot of ln(k3) versus 1000/T.
Figure 2. (A) Total catalyzed CO2 hydration rate kobs versus [OH] at 25 to 40 °C, with 6.25 mM total boric acid added. (B) Plot of k0 obs (= kobs k2[OH] k1) versus [B(OH)4] at 25 to 40 °C, with 6.25 mM total boric acid added. The addition of boric acid results in observed rates of CO2 hydration that are significantly above those found without boron addition. This increase in reaction rate (k0 obs) is linearly proportional to the tetrahydroxyborate ion concentration, [B(OH)4], which is proposed as the active form of boron. 4804
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Environmental Science & Technology stable intermediate that is rapidly hydrated. The individual reaction steps involved in this mechanism are listed below, and we see that the overall process achieves the conversion of CO2 into HCO3, along with the production of Hþ. It is thus apparent that the boric acid/borate system acts as a true catalyst for the CO2 þ H2O reaction. Interestingly, the mechanism revealed here for borate catalyzed CO2 hydration is the same as that commonly accepted for the enzyme carbonic anhydrase.8 In this mechanism the Zn2þ center in the neutral enzyme binds a water molecule, which then deprotonates at physiological pH (pKa ≈ 7) to form a Znþ(OH) site. Attack of CO2 on the OH moiety then produces HCO3 and returns the neutral enzyme form. Carbonic anhydrase is an incredibly efficient enzyme for CO2 hydration/dehydration, yet its thermal instability likely makes it unsuitable for solvent systems. The discovery that a small, stable inorganic molecule such as boric acid can undergo the same mechanism is significant, as it provides a framework for designing Scheme 1. Mechanism for the Borate-Catalyzed Hydration of CO2
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new catalysts that mimic the carbonic anhydrase mechanism, but are better suited to industrial CO2 capture applications. BðOHÞ3 3 H2 O T BðOHÞ4 þ Hþ BðOHÞ4 þ CO2 f BðOHÞ4 CO2 BðOHÞ4 CO2 þ H2 O f BðOHÞ3 3 H2 O þ HCO3 CO2 þ H2 O f HCO3 þ Hþ Figure 4A illustrates how the total CO2 hydration rate constant (kobs) varies as a function of both temperature and pH with 10 mM total added boric acid, based upon individual rate constant expressions measured in this study. Figure 4B demonstrates the percentage of kobs that can be attributed to the borate mechanism as a function of temperature and pH. At around pH 12 and above the CO2 hydration process is controlled almost exclusively by the CO2 þ OH reaction; however, it is difficult to maintain solvents at such high pH values as the chemical absorption of CO2 releases protons and acidifies the solution. For example, concentrated potassium carbonate solutions exist at around pH 11.5, but even with a small degree of CO2 loading the OH concentration drops significantly,19 providing a pH of 10 or below. This pH is typical of carbonate solvents used in gas absorption processes, which operate with partial loading of the solvent. It is thus apparent that under conditions relevant to solvent absorption the addition of borate compounds can provide a significant enhancement to the reaction rate. At a boron concentration of 0.65 M, typical of that used in carbonate absorption systems,13,14 the borate mechanism is the dominant CO2 hydration process at all relevant pH values, and at high temperatures (ca. 6080 °C) the hydration rate is almost independent of pH between pH 9 and 11 (see Figure S2, Supporting Information). These results also highlight two significant aspects that must be considered when designing new catalysts for CO2 capture: not only should the compound undergo relatively rapid reaction with CO2, it must also be present in the active form across the relevant pH range. Note that this analysis assumes no formation of polyborates and ignores activity effects, which will be important in concentrated carbonate/borate solutions, impacting the rate of all three hydration mechanisms.
Figure 4. (A) Total CO2 hydration rate as a function of pH (413) and temperature (2080 °C) with 10 mM total added boric acid. (B) Contribution of B(OH)4 ion catalysis in A. The addition of boric acid provides significant reaction rate enhancement at around pH 810. At lower pH the boron is present almost exclusively as B(OH)3, and the CO2 þ H2O mechanism dominates. At higher pH the OH ion concentration becomes significant, and the CO2 þ OH mechanism dominates. 4805
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Environmental Science & Technology Finally, this work also has important implications for CO2 absorption by lakes and oceans. Bodies of water exchange CO2 with the troposphere, where the CO2 þ H2O reaction (and to a lesser extent CO2 þ OH) ultimately results in the production of carbonate and bicarbonate ions. Understanding CO2 exchange between the atmosphere and the oceans is important in describing ocean acidification and carbon uptake; exchange to lakes plays a less important role in the global carbon cycle,20 but is critical to paleoclimate reconstructions and biogeochemical models. At high wind speeds CO2 uptake is controlled solely by mass transfer, but at low wind speeds the hydration of CO2 due to reaction with H2O and OH can become the controlling mechanism. This process results in some chemical enhancement of the gas absorption process, which is thought to provide a minor contribution to CO2 exchange with the oceans.21,22 Boron is a significant trace component of seawater, present at around 0.4 mM.23 Assuming an average surface water temperature of 15 °C and pH of 8.1,7 we predict that borate-catalyzed CO2 hydration accounts for 1.5% of the total hydration rate (Figure S3, Supporting Information) and is unlikely to impact upon CO2 exchange. Boric acid does however play an important role in sound absorption in seawater, due to relaxation to the borate form, and the chemistry identified here may help to explain the complex coupling observed between boric acid relaxation and the carbonate/bicarbonate system.24,25 Because the boron isotope ratio in carbonates has emerged as a useful proxy for ocean paleo-pH2628 it is critical that we understand all reactions in the boroninorganic carbon system (and their kinetics/equilibria).29,30 In alkaline saline lakes chemical enhancement can be more significant than for the oceans, increasing gas transfer rates by a factor of 2 or more; importantly, models tend to under-predict this rate of enhancement.22 Gas exchange models do not consider catalysis of CO2 hydration due to boron and other trace compounds, although potential catalysis by carbonic anhydrase was proposed15 and subsequently dismissed.3133 Boron (as well as other potential catalysts such as arsenic and phosphate) is found at high levels in certain saline lakes and the chemistry developed here should therefore be included in gas exchange models. For example, boron is present at around 30 mM in the highly saline (pH 9.7) waters of Mono Lake (California).34 At a typical temperature of 15 °C we estimate that borate catalysis would be the dominant CO2 hydration mechanism (Figure 5), increasing the reaction rate by 160% (Figure S4, Supporting Information). Mono Lake exhibits unique chemistry, with incredibly high levels of dissolved inorganic carbon and an elevated ratio of 14C to 12C that has proven difficult to explain.22,33,35,36 The sustained invasion of radiocarbon (14C) since the 1960s is particularly puzzling, as the driving force has been steadily dropping due to both increases in lake Δ14C and decreases in atmospheric bomb Δ14C (following the partial nuclear test ban treaty of 1963). This even led to the suggestion that radiocarbon was being illegally disposed of in the lake.33 The salinity of Mono Lake approximately doubled from the 1940s to 1980s, and we suggest that increasing boron concentrations over this time resulted in growing CO2 invasion rates, providing a new explanation for the carbon budget of Mono Lake. Our results might also help to explain the very high inorganic carbon content of other saline closed-basin lakes such as Borax Lake (California), Big Soda Lake, Pyramid Lake, and Walker Lake (Nevada), and Albert Lake (Oregon), which all have relatively high (J 1 mM) levels of boron (note that high levels of arsenic may also play a
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Figure 5. Predicted contribution from the water, hydroxide, and borate hydration mechanisms in Mono Lake water (30 mM total boron, 15 °C). The pH of Mono Lake is presently around 9.7, making the borate mechanism the main CO2 hydration pathway. Diversion of water from the lake resulted in an approximate doubling of salinity since the 1940s. A steadily increasing CO2 invasion rate from the 1940s to 1980s, due to increasing boron levels, provides a new explanation for the inorganic carbon content of Mono Lake.
role in Albert Lake).37 Each of these bodies of water has been found to exhibit significant chemical enhancement of CO2 gas transfer (Big Soda Lake, Mono Lake, Pyramid Lake),22 catalyzed CO2 hydration in laboratory experiments (Walker Lake),31 and/ or sustained carbonate precipitation (Big Soda Lake;38 Borax Lake, Albert Lake37). Borate catalysis may prove to be a general mechanism in saline lake chemistry that should be included in paleoclimate reconstructions and biogeochemical models.
’ ASSOCIATED CONTENT
bS
Supporting Information. Detailed description of rate constant calculations and CO2 þ H2O/OH experiments. Figures S1 to S7. Measured rate constants as a function of temperature.This material is available free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected]; phone: þ61 3 8344 6627; fax: þ61 3 8344 4153.
’ ACKNOWLEDGMENT We gratefully acknowledge financial support from the China Scholarship Council, the Cooperative Research Centre for Greenhouse Gas Technologies (CO2CRC), the Australian Government through its CRC Program, and the Particulate Fluids Processing Centre (PFPC). ’ REFERENCES (1) Lackner, K. S. Carbonate chemistry for sequestering fossil carbon. Annu. Rev. Energy Environ. 2002, 27, 193–232. (2) Lackner, K. S. A Guide to CO2 Sequestration. Science 2003, 300, 1677–1678. 4806
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