The Carbonate System in Natural Waters

Transcription

1 University of Miami Scholarly Repository Open Access Dissertations Electronic Theses and Dissertations The Carbonate System in Natural Waters Hector Bustos-Serrano University of Miami, Follow this and additional works at: Recommended Citation Bustos-Serrano, Hector, "The Carbonate System in Natural Waters" (2010). Open Access Dissertations This Open access is brought to you for free and open access by the Electronic Theses and Dissertations at Scholarly Repository. It has been accepted for inclusion in Open Access Dissertations by an authorized administrator of Scholarly Repository. For more information, please contact

2 UNIVERSITY OF MIAMI THE CARBONATE SYSTEM IN NATURAL WATERS By Héctor Bustos-Serrano A DISSERTATION Submitted to the Faculty of the University of Miami In partial fulfillment of the requirements for the degree of Doctor of Philosophy Coral Gables, Florida December 2010

3 2010 Héctor Bustos-Serrano All Rights Reserved

4 UNIVERSITY OF MIAMI A dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy THE CARBONATE SYSTEM IN NATURAL WATERS Héctor Bustos-Serrano Approved: Frank J. Millero, Ph.D. Professor of Marine and Atmospheric Chemistry Terri A. Scandura, Ph.D. Dean of the Graduate School Dennis A. Hansell, Ph.D. Professor of Marine and Atmospheric Chemistry Rana A. Fine, Ph.D. Professor of Marine and Atmospheric Chemistry Rik Wanninkhof, Ph.D. Adjunct Professor of Marine and Atmospheric Chemistry, Oceanographer at AOML NOAA

5 BUSTOS-SERRANO, HECTOR (Ph.D., Marine and Atmospheric Chemistry) The Carbonate System In Natural Waters (December 2010) Abstract of dissertation at the University of Miami. Dissertation supervised by Professor Frank J. Millero. No. of pages (130). Reliable measurements of the thermodynamics of the carbonate system are needed to better understand the CO 2 system in natural waters. New measurements of the carbonic acid pk 1 * and pk 2 * in seawater have been made over a wide range of temperatures (1 to 50 C) and salinities. The commonly used CO 2 constants of Mehrbach et al., (1973) were limited to salinities (19 to 43) and temperatures (2 to 35 C). They cannot be used to study estuarine or fresh waters. The results of these measured pk 1 * and pk 2 * values are in good agreement with those determined using the Miami Pitzer equations (Millero and Pierrot, 1998). The results in this dissertation demonstrate the validity of the model that can be used to study the carbonate system in most natural waters. The so called Miami model is presently being used to examine the effect of ocean acidification on natural waters The boric acid effect on the dissociation constants in seawater and NaCl solutions was tested. The addition of boric acid has little or no effect on pk 1 * values. However, the values of pk 2 *, decreases with the addition of small amounts of boric acid to ASW in agreement with the work of Mojica-Prieto and Millero (2002). The addition of larger concentrations of boric acid cause the values of pk 2 * to increase. These effects have been attributed to the interactions of boric acid with the carbonate ion (CO 2-3 ) in seawater (Mojica-Prieto and Millero, 2002).

6 The addition of boric acid to NaCl solutions in contrast, causes the values of pk 1 * and pk 2 * to decrease. This has been attributed to the interactions of borate ions with Mg 2+ and Ca 2+ in seawater. Further measurements in Na-Mg-Cl and Na-Ca-Cl solutions are needed to prove that this is the case. The boric acid effect on the carbonate constants indicate that an increase in boric acid has no affect on pk 1 *, but does change the values of pk 2 *. At low concentrations of boric acid, pk* 2 decreases, and at higher concentrations it increases. These results indicate that boric acid has some ionic interactions with the carbonate ion. Similar studies in NaCl indicate that both pk 1 * and pk 2 * decrease when boric acid is added. The differences between seawater and NaCl may be related to the interactions of Mg 2+ and Ca 2+ with borate anions. Further studies of NaCl with additions of MgCl 2 and CaCl 2 are needed to examine the effects in detail. Preliminary studies on the effect of DOC on the carbonate constants are not definitive. The change of the DOC concentration from 50 to 100 µmol kg -1 has little effect on the values of pk 1 * and pk 2 *. Dilutions of seawater with artificial seawater are complicated by changes in the concentration of boric acid. Earlier studies indicated that DOC may cause the 8 μmol kg -1 increase in total alkalinity of seawater needed to balance the thermodynamics of the system (Millero et al., 2002). This may be partially due to the new values for the B/Cl ratio in seawater found by Lee et al., (2010) that increases the TA by ~ 6 μmol kg -1. Further studies are needed to examine the effect of humic compounds in estuarine waters on the carbonate system. Measurements of ph or pco 2 along with TA and TCO 2 can be used to separate the effect of organic ligands on TA. If

7 DOC measurements are also made, one can relate the effect to organic ligands that can accept a proton. The cruises in the Little Bahama Banks show for the first time the active precipitation of CaCO 3 (Bustos-Serrano et al., 2009). This causes measured decreases of TA, TCO 2 and ph and increases in pco 2 in the whitings. This is in contrast to earlier studies on the Grand Bahama Banks where no active precipitation of CaCO 3 was every found (Morse et al., 2003; Millero et al., 2005). The differences appear to be due to the movement of fresh saturated seawater from the Gulf Stream into the LBB. The Gulf Stream water enters the GBB in the winter, and the precipitation occurs on the suspended sediment over the year. Observations are needed on the Grand Bahama Banks in the winter and throughout the year to prove that this is the case.

8 Dedication The way of God is good. His wonders are intended to be used by all, and exclusively by none. To God, all glory forever. To my dear parents, Graciela Serrano Avilés and Héctor Rubén Bustos Aldana, who gave me life, love and education. This work is dedicated to my forever kids Graciela del Carmen Bustos-Hernandez and Héctor Bustos-Hernández, in honor to their lives. To my wife María Esther Hernández Gámez who after all, is still my companion through this important moments in my life. To my brothers and sister, Francisco Rafael, José Rubén and Mariela, and their families, also my uncles, aunts, cousins and many friends who deserve special mention for walking together with me. To all my friends in México and because of those five years (+2 months) in Miami I must say all over the world. I will just mention your Country and you know that this dedication is especially for you: Spain, Italy, Venezuela, France, Australia, Tasmania, Germany, Chile, Costa Rica, Brazil, Portugal, Canada, Panamá, China, Colombia, Argentina, USA and of course Cuba. iii

9 Acknowledgments This work was possible thanks to the invaluable knowledge, patience, expertise, support and helpful advice of Dr. Frank J. Millero. Dr. Rana A. Fine, Dr. Dennis A. Hansell and Dr. Rik Wanninkhof. All of them deserve special mention for their helpful discussion of this dissertation. To my colleagues in the laboratory Valentina González- Caccia, Demetrio Milea, David Sergio Valdes-Lozano, Denis Pierrot, Yanxin Luo, J. Mike Trapp, Adriana Cabrera, Mareva Chanson, Xiaorong Zhu, Fen Huang, Vannessa Koehler, William T. Hiscock, Francisco J. Mojica-Prieto, Taylor B. Graham and Gay Ingram who deserves a special mention because of her unconditional friendship and kind attitude to review this and many other documents. The field work was possible thanks to the crew of the UM R/V Walton Smith. The funds for this work were provided by support from the Oceanographic Section of the National Science Foundation and the National Oceanic and Atmospheric Administration. iv

10 TABLE OF CONTENTS List of Figures List of Tables viii xiii Chapter 1. THE CARBONATE SYSTEM IN NATURAL WATERS Introduction and review of the carbonate system in seawater The Carbonate System Studies of carbonic acid dissociation constants in seawater Problems detected in the seawater values of pk 1 * and pk 2 *. 14 Chapter 2. EXPERIMENTAL METHODS FOR THE CARBONATE SYSTEM PARAMETERS General procedures Total alkalinity ph The Gran technique The determination of K 1 *K 2 * and K 2 * Fugacity of CO 2 (fco 2 ) Discussion of the pk 1 * + pk 2 * method Experimental methods and equipment for the pk 1 * and pk 2 * from titrations with hydrochloric acid Measurements of pk 1 * + pk 2 * from samples stripped of CO 2 and then with solid NaHCO 3 additions to a constant ph <steady state ph> Equipment and materials Reagents and solutions. 37 v

11 2.4 The effect of organic carbon on the carbonic acid dissociation constants Manipulation of seawater to gather information on carbonic acid dissociation constants Methods used in different media for carbonic acid pk i* Summary 45 Chapter 3. DISSOCIATION CONSTANTS OF CARBONIC ACID IN SEAWATER AS A FUNCTION OF SALINITY AND TEMPERATURE Modeling the carbonate system in natural waters Experimental methods and equipment Results and calculations Discussion Summary. 71 Chapter 4. EFFECT OF BORIC ACID AND DISSOLVED ORGANIC CARBON ON THE DISSOCIATION OF CARBONIC ACID Boric acid effect on the dissociation constants of carbonic acid Effect of DOC on the values of pk i * in seawater Results of carbonic acid pk i * in NaCl solutions Discussion Summary. 86 vi

12 Chapter 5. THE FORMATION OF WHITINGS ON THE LITTLE BAHAMA BANK Background Experimental methods General information Total alkalinity ph Total inorganic carbon Partial pressure of CO Calcium carbonate and water mass age Results Discussion Synopsis. 114 Chapter 6. SUMMARY AND CONCLUSION 117 Reference Section 120 vii

13 Figure 1.1. Figure 1.2. Figure 1.3. Figure 1.4. Figure 1.5. Figure 1.6. Figure 1.7. Figure 1.8. LIST OF FIGURES The increase of CO 2 in the atmosphere. Blue dots were measured on ice cores (Francey et al., 1999; Kawamura et al., 2003). Red dots are direct atmospheric measurements by D. Keeling at the Mauna Loa Observatory (Keeling and Whorf, 2006). 2 Increase in the DIC content and ph decrease in seawater as a function of the rise in atmospheric CO 2. (after Sabine et al., 2004a). 4 Scenarios of CO 2 in the atmosphere. (after Sabine et al., 2004a). 5 Comparisons of pk 1 * and pk 2 * in ASW Mehrbach et al., (1973) results in SW at S=35. The dotted lines represent 2σ error of the fits of the author s results as a function of temperature. They are in reasonable agreement (2σ = and 0.02; for pk 1 * and pk 2 * respectively). 11 Comparison of Δ(pK 2 *-pk 1 *) between Mehrbach et al., and different authors in ASW at S = 35. The dotted line represents 2σ error of the fits of the author s results to function of temperature. 12 Comparisons of pk 1 * and pk 2 * values between Mojica- Prieto and Millero (2002) and others at S = 35.0 as a function of temperature. The dotted lines represent 2σ errors of the fits. 13 Differences between measured and calculated values of fco 2 from an input of TA and TCO 2 as a function of fco 2 using the combined equations (equations 22 and 23 from Mojica-Prieto and Millero, 2002) which uses Mehrbach et al., (1973) and Mojica-Prieto (2001). Panels (A) and (B) were determined without and with correction for changes in pk* 2 as a function of TCO 2 (eq. 1.12). The solid line with slope in (A) is an estimate of the offset at high fco Differences between the measured and calculated values of fco 2 from an input of TA and TCO 2 as a function of TCO 2. Panel A) is using the combined equations of Mehrbach et al., (1973) and Mojica-Prieto (2001). Panel B) is using the correction changes in pk 2 * as a function of TCO 2 (eq. 1.12). 16 viii

14 Figure 2.1. Figure 2.2. Figure 2.3. Figure 2.4. Figure 3.1. Figure 3.2. Figure 3.3. Figure 3.4. Figure 3.5. Figure 3.6. Figure 3.7. Figure 3.8. Potentiometric determination of ½(pK 1 * + pk 2 *) using pure solid NaHCO 3, this is the steady state ph or ph Closed cell for the titration of NaCl solutions, ASW and SW with HCl ~0.25M. 35 Typical titration for Total Alkalinity using HCl ( 0.25 N). A high resolution curve was setup for detail carbonic acid pk i * fittings in all NaCl solutions. 36 Typical titration for Total Alkalinity in SW and ASW with Na 2 CO 3 added. A detailed high resolution curve is obtained from the titration setup to better fit the carbonic acid pk i *. 37 A comparison of the values of pk 1 * and pk 2 * from 0 to 40 C and S = 35 determined by various workers with the model of Millero and Roy (1997). 48 A comparison of the values of pk 1 * and pk 2 * at 25 C as a function of the square root of salinity (S = 0 to 45) of Mehrbach et al., (1973) with the model results from Millero and Roy (1997). 51 The measured values of pk 1 * and pk 2 * from 10 to 50 C as a function of the square root of ionic strength (I). The smooth curves are the values calculated from the model of Millero and Roy (1997). 60 The differences between the measured and fitted values of pk 1 * as a function of salinity and temperature. 62 The differences between the measured and fitted values of pk 2 * as a function of salinity and temperature. 63 A comparison of our pk 1 * and pk 2 * results at S = 35 with literature values. 65 Comparison of this work for pk 1 * and literature values as a function of temperature and salinity. The standard error of the fit for all the measurements is shown (see Table 3.4). 66 The differences between all the measured and fitted values of pk 2 * as a function of temperature and salinity. 68 ix

15 Figure 3.9. Figure Figure 4.1. Figure 4.2. Figure 4.3 Figure 4.4. Figure 4.5. Comparison of the new equation for pk 1 * (Millero et al., 2006) with earlier workers at Salinity = 35. Mehrbach et al., (1973); Hansson, (1973); Goyet and Poisson, (1989); Roy et al., (1993a); Mojica-Prieto and Millero, (2002). 71 Comparison of the new equation for pk 2 * (Millero et al., 2006) with earlier workers at Salinity = 35. Artificial Seawater measurements of Hansson, (1973); Goyet and Poisson, (1989) and Roy et al., (1993a). Seawater measurements of Mehrbach et al., (1973) and other works by Dickson and Millero (1987). 72 Effect of boric acid (HB) on real seawater carbonic acid pk 2 *. The red dot is the expected pk 2 * (8.99) when HB = 0. The regression is pk 2 * = E -4 * HB E -7 * HB The difference in the value of pk 2 * with added boric acid in seawater compared to the pk 2 * values when HB=0 from Figure DOC effect on the carbonic acid pk 2 * value. Solid dots are diluted seawater from the Gulf Stream. pk 2 * Open circles are natural seawater from Gulf Stream or Biscayne Bay (Measured Millero et al., 2006). 79 Values of pk 1 * in sodium chloride 0.7m solution with Na 2 CO µm at 25 C over different concentrations of boric acid. The estimated standard error of the pk 1 * value is and for HB=0 the estimated pk 1 * value is Values of pk 1 * on sodium chloride 0.7m solution with Na 2 CO µm at 25 C over different loads of boric acid. The standard error of for the pk 1 * estimate is very good, considering this large amount of HB. 82 Figure 4.6. Values of pk 2 * on sodium chloride 0.7m solution with 2000 µm Na 2 CO 3 at 25 C over different loads of boric acid. The standard error of the estimate is for pk 2. The pk 2 * = when no boric acid is added. 83 Figure 4.7. Temperature effect on carbonic acid pk* 1 in 0.25 m NaCl solution with added carbonate (250 µm) and HB (250 µm). The standard error of the estimate is and the pk 1 *(0) = x

16 Figure 4.8. Figure 5.1. Figure 5.2. Figure 5.3. Figure 5.4. Figure 5.5. Figure 5.6. Temperature effect on carbonic acid dissociation pk 2 * in 0.25m NaCl solution with 250 µm Na 2 CO 3 and 250 µm boric acid. The standard error of the estimate is The pk 2 *(0) estimate at T = 0 C is The maximum CaCO 3 (s) precipitated on the Grand Bahama Bank over the time of the year. The values are based on the changes in the total alkalinity corrected for salinity. 89 The maximum CaCO 3 (s) precipitated from waters on the Grand Bahama Bank as a function of the salinity. 90 A photograph (Roll 719; Frame 29) of whitings on the Little Bahama Bank from the NASA Space Shuttle. Photograph courtesy of the Image Science & Analysis Laboratory, NASA Johnson Space Center ( 93 Contours of the surface salinity of waters in the Little Bahama Bank from cruise in July 2003 and May Blue circles are sampling stations. Interpolation using DIVA from Ocean Data View software. 96 Contours of the normalized total alkalinity of waters in the Little Bahama Bank from cruises in July 2003 and May Contours of the normalized total carbon dioxide of waters in the Little Bahama Bank from cruises in July 2003 and May Figure 5.7. ph contours in the Little Bahama Bank from cruises in 2003 and Figure 5.8. Figure 5.9. Figure Figure Figure Figure Values of NTA plotted against salinity which is a proxy for water residence time with the highest salinity occurring at about 144 d. 101 Contours of the saturation state of aragonite of waters in the Little Bahama Bank from cruises in 2003 and The values of NTA plotted for the stations occupied on the cruises. 107 The values of NTCO 2 plotted for the stations occupied on the cruises. 108 The values of pco 2 plotted for the stations occupied on the cruises. Reference line at 350 µatm is the atmospheric value. 109 The amount of suspended CaCO 3 (s) in whiting and non whiting waters. Note the difference in the scales between open ocean and whitings waters. 110 Figure The Sr/Ca versus the Mg/Ca ratios of whiting and fine- 111 xi

17 Figure Figure Figure Figure grained bottom LBB sediment. Whiting values are solid circles and bottom sediments are open circles. δ 13 C versus δ 18 O, relative to the PDB standard, of whiting and bottom sediments. Whiting values are solid circles and bottom sediments are open circles. 111 Little Bahama Bank (pale blue) just north of the Great Bahama Island show some whitings using satellite imagery on Google Earth (Oct. 25, 2008). 114 Cruise Track at LBB cruise in May 2005 (blue line) some whitings were just observed from the R/V Walton Smith bridge. Blue dots reflect the size of whitings, most of them with irregular shapes. 115 Whiting observed at the LBB cruise in July 2003 on the bridge of the R/V Walton Smith, University of Miami. 116 xii

18 LIST OF TABLES Table 1.1. Table 1.2. Sources and sinks of CO 2. Oceanic removal is an important sink of carbon. Modified after Le Quere et al., (2009) and Sabine and Tanhua, (2010). 3 Summary of measurements made on the dissociation constants of carbonic acid in SW and ASW by various workers at S=35 and t=25 C compared with the computer code CO 2 Brine.bas, numbers in parenthesis indicate the standard error, all units in SW scale (mol Kg -1 soln.). 7 Table 2.1. Composition of artificial seawater in moles kg -1 solution used by various authors. 41 Table 3.1. Table 3.2. Table 3.3. Table 3.4. Table 3.5. The effects of TA (µmol kg -1 ) and TCO 2 (µmol kg -1 ) levels and ph on the determination of the pk* 1 and pk* 2 at various temperatures. 55 Measured values of pk 1 * and pk 2 * for carbonic acid in seawater as a function of salinity and temperature. 56 Coefficients for the fits of the values of pk* 1 and pk* 2 in seawater as a function of temperature, salinity and ionic strength. 64 Coefficients for the fits of the values of pk* 1 and pk* 2 in seawater as a function of temperature and salinity (measurements combined with literature studies). 67 Comparisons of the standard deviations of the authors and the standard error of the fit for all the measurements. 69 xiii

19 Table 3.6a. Table 3.6b. Table 4.1. Table 4.2 Table 4.3. Table 5.1. Table 5.2. Errors in the determination of CO 2 parameters at S = 35 and t = 25 C due to errors of in pk 1 *. 70 Errors in the determination of CO 2 parameters at S = 35 and t = 25 C due to errors of in pk 2 *. 70 Results of the boric acid additions to Gulf Stream seawater (S=36.101) carbonic acid dissociation constants at t = 25 C 77 Dissociation constants of carbonic acid in mixtures of Gulf Stream seawater and ASW with different loads of DOC and HB (TA and TCO 2 in µmol kg -1 ). Salinity constant at Diluted ASW results on steady state ph 0, ½(pK 1 +pk 2 ) and carbonic acid stoichiometric dissociation constants. Included the results of ph 0 for Mojica-Prieto and Millero (2002) = M&M. 87 Carbonate measurements made on July 2003 Bahamas Cruise. 103 Carbonate measurements made on May 2005 Bahamas Cruise. 105 xiv

20 Chapter 1 THE CARBONATE SYSTEM IN NATURAL WATERS 1.1 Introduction and review of the carbonate system in seawater. The carbonate system is part of the overall biogeochemical carbon cycle in natural waters. It provides the acid-base buffering of the oceans, the carbon needed for primary production and controls the solubility and precipitation of calcium carbonates. The fundamental thermodynamic equilibria involved are divided into the solubility of carbon dioxide, distribution of carbon dioxide between seawater and air, hydration of carbon dioxide, the dissociation of carbonic acid and the solubility of CaCO 3 (s) minerals The Carbonate System. Interest in changes in the distribution of CO 2 in the oceans is related to the need to understand how the increase of atmospheric CO 2 and the increase in temperature will affect the climate. It is well documented that the CO 2 levels in the atmosphere are increasing (from 280 ppm in preindustrial times to more than 388 ppm at the present time) due to the burning of fossil fuel (Keeling et al., 1995; Keeling and Whorf, 2002; Somova et al., 2003). Continuous CO 2 atmospheric measurements made since 1958 at the Mauna Loa Observatory in Hawaii demonstrate the secular increase in atmospheric CO 2 (Fig. 1.1). These results have been supplemented with preindustrial levels by the CO 2 trapped in ice (Francey et al., 1999; Kawamura et al., 2003). These studies demonstrate that the levels of CO 2 in the atmosphere have never been above 300 ppm over the last 800,000 years (Francey et al., 1999). 1

21 2 Figure 1.1. The increase of CO 2 in the atmosphere. Blue dots were measured on ice cores (Francey et al., 1999; Kawamura et al., 2003). Red dots are direct atmospheric measurements by D. Keeling at the Mauna Loa Observatory (Keeling and Whorf, 2006). CO 2 is a greenhouse gas, meaning that it absorbs solar radiation transmitted back from Earth in the infrared. This energy is partially dissipated as heat, warming the atmosphere, land and sea. Terrestrial plants play a major role taking up CO 2 from the atmosphere (e.g. Falkowski, 2002). The oceans remove close to 40% of anthropogenic CO 2 (Table 1.1.) mainly by the solubility and biological pumps (e.g. Falkowski, 2002). The cold high latitude waters take up most of the CO 2, and it is stored in sinking deep waters for hundreds of years. The uptake of CO 2 by the ocean has been monitored as a function of time by a number of programs (e.g. Joint Global Ocean Flux Studies (JGOFS), World Ocean Circulation Experiment (WOCE), Ocean Atmosphere Carbon

22 3 Exchange Program (OACES), and the Climate Variability and Predictability Research Program (CLIVAR). Table 1.1. Sources and sinks of CO 2. Oceanic removal is an important sink of carbon. Modified after Le Quere et al., (2009) and Sabine and Tanhua, (2010). Average Perturbation (Pg C yr -1 or10 15 g C yr -1 ) Sources Fossil Fuel Combustion 8.7±0.5 Deforestation 1.4±1.0 Total 10.1±1.4 Sinks Atmosphere 4.1±0.2 Oceans N. Hemisphere forest grow 0.4±0.4 Other terrestrial sinks 2.7±1.0 Total 10.5±1.5 Evidence of the CO 2 increase in the oceans and its impact has been presented (Brewer, 1978; Chen and Millero, 1979; Siegenthaler and Sarmiento, 1993; Gruber et al., 2001; Orr et al., 2001; Wallace, 2001; Feely et al., 2004; Sabine et al., 2004a; Orr et al., 2005). More CO 2 measurements are in progress in order to understand the mechanism and rate of the uptake of fossil fuel by the oceans and land (Libes, 1992; Quay et al., 1992; Goyet and Brewer, 1993; Tsunogai et al., 1993; Gruber et al., 1996; Goyet et al., 1998; Peng et al., 1998; Sabine et al., 2004b; Beaulieu et al., 2010; Peng and Wanninkhof, 2010). The distribution of the components of the carbonate system (carbonic acid, H 2 CO 3 ; bicarbonate, HCO 3 - and carbonate, CO 3 2- ) in the oceans is shown in Fig. 1.2 as the CO 2 increases in the atmosphere CO 2 (Sabine et al., 2004a). The change in the

23 4 concentration of CO 3 2- and lowering of the ph will affect the precipitation, the dissolution of carbonate minerals, and the speciation of metals in seawater. Figure 1.2. Increase in the DIC content and ph decrease in seawater as a function of the rise in atmospheric CO 2. (after Sabine et al., 2004a). The possible scenarios for the future increase of CO 2 in the atmosphere due to the burning of fossil fuel are shown in Fig These scenarios are based on estimates for the burning of fossil fuels (Sabine et al., 2004a). Some predictions using these scenarios on the marine productivity have been proposed on different ecosystem structure and complexity, with a decrease in the global mean productivity from 2% to 20% by the year 2100 relative to preindustrial concentrations (Steinacher et al., 2010).

24 5 Figure 1.3. Scenarios of CO 2 in the atmosphere. (after Sabine et al., 2004a). Once the atmospheric CO 2 enters the ocean surface water, it is hydrated to form carbonic acid (H 2 CO 3 ), which dissociates as a dibasic acid governed by the equilibrium constants for the dissolution of CO 2 and dissociation of H 2 CO 3 and HCO - 3 (Fig. 1.2). The thermodynamic equilibrium constants are given by K 1 = a(h + ) a(hco 3 - ) / a( H 2 CO 3 ) (1.1) and the K 2 = a(h + ) a(co 3 2- ) / a(hco 3 - ) (1.2) These thermodynamic equilibrium constants are expressed in terms of activities of the ions a(i) involved and vary with temperature, ionic strength and the ionic medium. The

25 6 latter exerts its effects by changing the activities of the ions and thus the ratios of their concentrations one to another. Studies of the CO 2 system in seawater use stoichiometric constants (K i*) to examine the equilibria. CO 2 (g) = CO 2 (aq) K 0 * (1.3) CO 2 + H 2 O = H + + HCO 3 - K 1 * (1.4) HCO 3 - = H + + CO 3 2- K 2 * (1.5) These stoichiometric constants are defined in terms of the concentration of the species [i] in mol kg -1 K 0 * = [CO 2 ]/pco 2 (1.6) K 1 * = [H + ] [HCO 3 - ]/a{h 2 O} [CO 2 ] (1.7) K 2 * = [H + ] [CO 3 2- ]/[HCO 3 - ] (1.8) These stoichiometric constants are related to the thermodynamic values by K 0 = K 0 * γ(co 2 ) (1.9) K 1 = K 1 * γ(h+) γ(hco 3 - )/γ(co 2 ) a{h 2 O} (1.10) K 2 = K 2 * γ(h + ) γ(co 3 2- )/γ(hco 3 - ) (1.11) where γ(i) is the activity coefficient of species i and pco 2 is the vapor pressure of CO 2. The activity coefficient (γ i ) for the individual species can be calculated using an ionic

26 7 interaction model (Millero, 1982; Millero and Pierrot, 1998). The measured values are in good agreement with Pitzer ionic interaction models (Millero and Roy, 1997; Millero and Pierrot, 1998; Millero et al., 2006). Laboratory measurements of the values of pk i * as a function of salinity, temperature and pressure (Millero, 1995; Millero et al., 2006) have been fitted to equations that can be used to determine the components of the CO 2 system in the oceans. A summary of the measurements of the carbonate constants is given in Table 1.2. Table 1.2. Summary of measurements made on the dissociation constants of carbonic acid in SW and ASW by various workers at S = 35 and t = 25 C compared with the computer code CO 2 Brine.bas. Numbers in parentheses indicate the standard error. All units in SW scale (mol kg -1 soln.). Thermodynamic data for Carbonic Acid Temp = 25ºC; Sal = 35.0 Author Year t ºC Salinity * pk 1 * pk 2 Hansson (ASW) (±0.007) (±0.009) Mehrbach et al., (SW) (±0.006) (±0.010) Goyet and Poisson (ASW) (±0.007) (±0.011) Roy et al., (ASW) (±0.002) (±0.003) Mojica-Prieto and Millero (SW & ASW) (±0.003) (±0.004) Millero et al., (SW) to 50 1 to (±0.013) (±0.020) CO 2 Brine.bas

27 8 1.2 Studies of carbonic acid dissociation constants in seawater. The CO 2 system in the ocean can be studied by measuring at least two of the following parameters: total CO 2 (TCO 2 ); total alkalinity (TA); ph and fugacity of CO 2 (fco 2 ). The internal consistency of the field data can be checked by measuring all four parameters. Computer software (e.g. CO2SYS) has been developed to use the stoichiometric constants of carbonic acid with any two parameters allows calculation of the other parameters. The carbonate constants of Millero et al. (2006) are available in the recent Excel version of CO2SYS from the CDIAC website ( These constants should be used for studies of seawater as a function of salinity from 0 to 50 and temperature from 0 to 50 o C. The constants of Mehrbach et al. (1973) are not valid below S = 20. The computer program with an input of any two CO 2 parameters produces estimates of the other two parameters as well as the components of the CO 2 and CaCO 3 system. Over the last 10 years, significant improvements have been made on the measuring of these four parameters (D.O.E., 1991; 1994, Dickson et al., 2007); a CO2SYS electronic Excel version is available from CDIAC ( The use of certified reference materials (CRM) developed by Dr. Andrew Dickson (Scripps Institution of Oceanography) a number of years ago has improved the precision and accuracy of TA and TCO 2 measurements at sea. Improvements in the techniques for measuring the CO 2 system in the ocean have motivated the re-examination of the internal consistency of the CO 2 parameters in the laboratory (Lee and Millero, 1994; Lee et al., 1996; Lee et al., 1999; Lueker et al., 2000)

28 9 and at sea (Clayton and Byrne, 1993; Millero et al., 1993b; Lee et al., 1997; Byrne et al., 1999; Wanninkhof et al., 1999; Millero et al., 2002). However, there are some discrepancies that exist and these will be addressed in this thesis. The addition of boric acid to artificial seawater (ASW) can account for some of the differences of K 2 * in real (SW) and ASW seawaters (Mojica-Prieto and Millero, 2002). The causes of these differences are still not well known. Some workers (Lee et al., 1996; Lueker et al., 2000; Millero et al., 2002) mention that there is also a TCO 2 effect on the pk 1 * and pk 2 * values. This effect on pk 2 * can be related to intrinsic effects of CO 2 levels on the thermodynamics of the carbonate system, type of organic material in seawater (regardless the low concentrations) and more likely, carbonate-borate interactions. The effect on the pk 2 * value can lead to errors calculating fco 2 from TCO 2 and TA. One of the earliest records on boron in seawater and in organisms is from Igelsrud et al., (1938). The new B/Cl measurements of Lee et al., (2010) affect TA of surface waters by as much as 6 μmol kg -1. This may be part of the cause of pk 2 * being different in SW and ASW. Using the B/Cl ratios between Uppstrom (1974) and Lee et al., (2010) and the thermodynamic calculations, a difference of 5 µmol kg -1 is found when converting TA to CA and vice versa. The most reliable earlier carbonic acid dissociation constants for seawater (SW) were determined by Mehrbach et al., (1973). The comparisons are shown in Fig. 1.4 for the values of pk 1 * and pk 2 * in ASW and SW as a function of temperature and salinity. The differences in pk i * are within and 0.04 respectively for pk 1 * and pk 2 * and the

29 10 measurements by Mojica-Prieto and Millero (2002). These constants are valid for field (Millero et al., 2002) and laboratory measurements (Lee and Millero, 1994; Lee et al., 1996; Lueker et al., 2000) in real seawater, but are not valid below S = 20. Since the values of pk 1 * from the studies in SW and ASW agree within 0.01, the differences in pk 2 * - pk 1 * shown in Fig. 1.5 are largely due to uncertainties in pk 2 *. A reliable value of pk 2 * - pk 1 * is required when calculating fco 2 using TA and TCO 2. The carbonateborate interactions may cause the differences in the pk 1 * and pk 2 * values as suggested by Mojica-Prieto and Millero, (2002) between ASW and SW. Since the pk 1 * value is easier to determine, I have determined values of pk 2 * by measuring the ph = ½ (pk 1 * + pk 2 *), as developed by Mehrbach et al., (1973) and improved by Mojica-Prieto and Millero (2002) with an electrode calibration to measure ph and the use of the spectrophotometric ph technique of Clayton and Byrne (1993). The precision and accuracy of the modern equipment and techniques provide an opportunity to refine the thermodynamic models. The present analytical precision in the measurements of the carbon dioxide system parameters is ± for ph, ± 1 μmol kg -1 soln. for TA, ± 1 μmol kg -1 soln. for TCO 2 and ± 1 μatm for fco 2 (Clayton and Byrne, 1993; Johnson et al., 1993; Wanninkhof and Thoning, 1993; Yao and Byrne, 1998; Millero et al., 2002; Dickson et al., 2003).

30 ASW - Mehrbach et al., (S = 35) 0.02 pk* pk* Hansson Goyet and Poisson Roy et al Temperature ( C) Figure 1.4. Comparisons of pk 1 * and pk 2 * in ASW Mehrbach et al., (1973) results in SW at S=35. The dotted lines represent 2σ error of the fits of the author s results as a function of temperature. They are in reasonable agreement (2σ=0.014 and 0.02; for pk 1 * and pk 2 * respectively).

31 ASW - Mehrbach et al., (S = 35) Hansson Goyet & Poisson Roy et al., (pk 2 - pk 1 ) σ Temperature ( o C) Figure 1.5. Comparison of Δ(pK 2 *-pk 1 *) between Mehrbach et al., and different authors in ASW at S = 35. The dotted line represents 2σ error of the fits of the author s results to function of temperature. Mojica-Prieto and Millero, (2002) and Millero (2007) pointed out that over decades of studies there are no significant differences between pk 1 * values that lie between ± 0.01 (Fig. 1.6). However, that is not the case for pk 2 * with higher differences of ± 0.04; such discrepancies in the carbonic acid dissociation constants can account for larger errors when this uncertainty is introduced in the analytical procedures of the carbonate system. In particular, high deviations are observed when estimating TA and TCO 2 from ph-fco 2, and also on the estimation of fco 2 from TA and TCO 2 data.

32 pk 1 * Hansson Goyet and Poisson Roy et al. Mojica-Prieto and Millero pk 2 * Temperature ( C) Figure 1.6. Comparisons of pk 1 * and pk 2 * values between Mojica-Prieto and Millero (2002) and others at S = 35.0 as a function of temperature. The dotted lines represent 2σ errors of the fits.

33 Problems detected in the seawater values of pk 1 * and pk 2 *. The carbonic acid pk s are affected by many factors, temperature, pressure, concentration and presence of interferences (Millero, 1995; Dickson, 2004a; Millero et al., 2006). As a result of years of investigations, the University of Miami s, Marine Physicochemical Group published the stoichiometric constants for carbonic acid in seawater (Millero et al., 2006); then, a year later, the same group published the pk i * for the same acid in sodium chloride solutions (Millero et al., 2007). Still, a large difference of 0.04 for pk 2 * prevails between SW and ASW. Thus, some aspects need to be considered: There is a difference in values of pk 2 * between SW and ASW (Fig. 1.4) that appears to be due to differences in the composition of the two solutions (Mojica-Prieto and Millero, 2002). This difference is related to changes in the activity coefficients of CO 2-3 in the different media. Applying internal consistency tests to the system (Lueker et al., 2000; Lee et al., 2000); showed that there is an influence of pk 2 * on fco 2 that agrees with results from the field measurements shown in Figure 1.7 (Millero et al., 2002).

34 15 fco 2 (meas - calc) A fco 2 (µatm) fco 2 (meas - calc) B fco 2 (µatm) Figure 1.7. Differences between measured and calculated values of fco 2 from an input of TA and TCO 2 as a function of fco 2 using the combined equations (equations 22 and 23 from Mojica-Prieto and Millero, 2002) which uses Mehrbach et al., (1973) and Mojica-Prieto (2001). Panels (A) and (B) were determined without and with correction for changes in pk 2 * as a function of TCO 2 (eq. 1.12). The solid line with slope in (A) is an estimate of the offset at high fco 2. Lee et al. (2000) found that an offset occurs at high fco 2 for some of the cruises where fco 2 was measured by the GC-FID system (Neill et al., 1997). The average differences in fco 2 are 15.8 µatm while the standard errors are 23.3 µatm. The deviations at fco 2 above ~600 µatm are greater than at low fco 2, apparently due to changes in the pk 2 * - pk 1 * (Lee et al., 2000; Lueker et al, 2000). The differences between the measured and calculated values of fco 2 as a function of TCO 2 (Fig. 1.8) show larger deviations at high TCO 2.

35 fco 2 (meas - calc) A TCO 2 (µmol kg -1 ) 150 fco 2 (meas - calc) B TCO 2 (µmol kg -1 ) Figure 1.8. Differences between the measured and calculated values of fco 2 from an input of TA and TCO 2 as a function of TCO 2. Panel A) is using the combined equations of Mehrbach et al., (1973) and Mojica-Prieto (2001). Panel B) is using the correction for changes in pk 2 * as a function of TCO 2 (eq. 1.12). At values of TCO 2 < 2050 µmol kg -1, the deviations are independent of the TCO 2. By adjusting the values of pk 2 * above 2050 µmol kg -1, it was possible to lower the average deviations to zero and the standard error was 19.3 µatm (see Figures 1.7B and 1.8B). This resulted in the relationship (Millero et al., 2002): pk 2 * TCO2 = pk 2 * x10-4 (TCO ) (1.12) which is valid at 20 o C and at TCO 2 > 2050 µmol kg -1.

36 17 The studies of Mojica-Prieto and Millero (2002) indicate that the differences in the values of pk 1 * and pk 2 * in real and artificial seawater are related to interactions of the borate and carbonate systems. At 25 o C the values of pk 1 * in artificial seawater with boric acid are ~0.01 lower than in artificial seawater without boric acid; while, the values of pk 2 * in artificial seawater with boric acid are higher ~0.04 than seawater without boric acid (Mojica-Prieto and Millero, 2002). The interactions responsible for these differences are not clear. The small increase in K 1 * in SW can be attributed to a decrease in the activity coefficient of HCO - 3 (γ HCO3 ) or to experimental error (Millero et al., 2007). 2- The decrease in K 2 * can be attributed to an increase in the activity coefficient of CO 3 (γ CO3 ). Assuming that the B(OH) - 4 repulsive interactions between HCO - 3 and CO 2-3 are small, the changes in γ CO3 can be attributed to interactions with B(OH) 3. McElligot and Byrne (1998) have shown the interactions of B(OH) 3 and HCO 3 - can be due to the formation of a borate carbonate complex: HCO B(OH) 3 = B(OH) 2 CO H 2 O (1.13) The formation of B(OH) 2 CO - 3 however is not large enough to significantly affect the value of pk 1 * in SW. Similar interactions between B(OH) 3 and CO 2-3 would not be expected to increase γ CO3. Further studies of boric and carbonic acid mixtures are needed to elucidate these interactions (Millero et al., 2002; Mojica-Prieto and Millero, 2002; Millero et al., 2006) as well as the possible ~6-8 µmol kg -1 effect on TA from organic ligands (Bradshaw and Brewer, 1988; Millero et al., 2002). This is a part of my research and is covered in Chapters 2 and 3. Chapter 2 discusses the methods to determine organic carbon and the carbonate parameters for this study made in the laboratory and at sea.

37 18 Chapter 3 discusses new measurements of the pk 1 * and pk 2 * of seawater from 0 to 50 o C and S = 0 to 50. These measurements confirm the results of Mehrbach et al. (1973) and Mojica-Prieto and Millero (2002). They also provide carbonate constants that can be used in estuarine waters. The earlier studies were made for salinities >20. Chapter 4 discusses the results of measurements of stoichiometric constants for the dissociation of carbonic acid in NaCl solutions, SW and ASW with and without boric acid as a function of ionic strength and temperature. Measurements were also made to examine how dissolved organic carbon (DOC) may affect the carbonic acid constants. Chapter 5 describes the results of two Little Bahama Bank cruises on board the UM R/V Walton Smith (July 2003 and May 2005) that studied whitings (patches of suspended fine-grained calcium carbonate). These whitings have been studied on the Great Bahama Bank (GBB) since the 1940s. The source and cause of these whitings have been hotly debated for a number of years. Recent studies have shown that resuspension of underlying sediments act as seed crystals for the slow precipitation of CaCO 3. Similar studies have not been conducted on the nearby, more northerly Little Bahama Bank (LBB) where satellite photographs indicate the presence of extensive whitings north of the Grand Bahama Island. Research cruises were made to the LBB in July 2003 and May On board measurements were used to examine the distribution of the components of the carbonic acid system in LBB waters. The distribution of suspended calcium carbonate was also determined and samples of the suspended material and underlying sediments were collected for analyses, including 14 C age determinations. Chapter 6 summarizes the results and conclusions of the studies made for this dissertation.

38 Chapter 2 EXPERIMENTAL METHODS FOR THE CARBONATE SYSTEM PARAMETERS 2.1 General procedures. Experimental methods are presented for the following carbonate parameters: Total Alkalinity (TA), ph, the Gran technique, fugacity of CO 2 (fco 2 ). These methods are used in the following chapters Total alkalinity. TA was determined by adding HCl past the carbonic acid end point using an automated potentiometric titration system (Millero et al., 1993a). The Nernstian behavior of the electrodes was checked by titrating pure 0.7 M NaCl solutions with a high resolution titration setup and processing the e.m.f. and volume of HCl data with the ESAB2M Qbasic program (DeStefano et al., 1987) developed at the University of Messina, Italy. This program provides a good estimation of the standard potential of the cell (E*). Each titration is calibrated with its own E* value, which is adjusted by nonlinear least squares procedure (Dickson et al., 2007) ph. The initial ph was measured from the sample titration using the first point (no HCl addition) on the seawater scale ph (SWS) : [H + ] = H + + HSO HF. The ph is also evaluated on this scale using a spectrophotometer for artificial seawater (ASW) and real seawater (SW), where a sample is transferred to a closed 10 centimeter Quartz optical cell with the spectrophotometric technique described by Clayton and Byrne (1993) using m-cresol purple as indicator. The addition of 20 μl of an 8 mm solution of the indicator with a Gillmont pipette yielded a final concentration of the indicator in the sample of 19

39 20 approximately 5.3x10-6 M. The absorbance measurements were performed with a Hewlett Packard 8453 UV-Visible diode array, single beam, and microprocessor-controlled spectrophotometer with collimating optics. During measurements, the optical cells were placed in a thermally controlled chamber that was connected to the NESLAB circulating temperature bath. The spectrophotometric measurements were made to increase the precision and accuracy of the ph determined by potentiometry in the ASW and SW samples. The spectrophotometric technique has a precision of ± , which is nearly an order of magnitude better than the potentiometric determination (± 0.002). For the ½(pK 1 +pk 2 ) procedure (described later in this section) a sample is measured with the ph spectrophotometric technique from the titration alkalinity cell when the system is in the steady state ph (ph 0 ). The equation used to estimate the ph from the absorbance measurements is (Clayton and Byrne, 1993): ph total = pk indicator + log [ ( R ) / ( R ) ] (2.1) In equation (2.1), R is the ratio of the indicator absorbance (A) at molar absorptivity maxima and can be calculated as: R = (A λ=578nm - A λ=730nm ) / ( A λ=434nm - A λ=730nm ) (2.2) The dissociation constant for the m-cresol purple can be calculated from the equation (Clayton and Byrne, 1993): pk indicator = / T ( 35 S ) (2.3) where T is the absolute temperature (from to K) and S is the salinity (from S= 30 to 37) of the sample. The values of the pk indicator and ph are on the total proton scale in units of mol kg -1. The calculated ph total has to be corrected for a ph change

40 21 resulting from addition of the dye. The correction is valid for each batch of indicator solution. The equation used to correct the ph is: ph total corrected = ph total ( 7.9 ph total ) / 150 (2.4) Lee and Millero (1995) re-determined the value of pk indicator at S = 35 from 0 to 40 C with a 0.04 M TRIS buffer. The ph of the buffer was determined with an electrode of the type H 2, Pt AgCl, Ag. They fitted their results to the following equation: pk indicator = / T log T (2.5) where T is the absolute temperature. The dissociation constant is on the total proton scale and has units of molality. The use of equations 2.1 to 2.7 assumes that molar absorptivity ratios are independent of temperature. The equation of (Lee and Millero, 1995) in combination with the salinity dependence term (Clayton and Byrne, 1993) is used to estimate the spectrophotometric ph. The final equation is given by (mol kg -1 soln): pk indicator = /T logt (35 S) + log( *s) (2.6) Transformation to the seawater scale is made using Dickson, (1984) and Millero, (1996): ph sws = ph total log [ (1 + S T / K HSO4- + F T / K HF ) / (1 + S T / K HSO4- ) ] (2.7) where S T is the total concentration of sulfate ion and F T is the total concentration of fluoride ion in the seawater sample. As part of the analytical routine, CRM s were often used for the quality control of oceanic carbon dioxide measurements, as suggested by Dickson et al. (2003).

41 The Gran technique. This technique is used to estimate the first and second dissociation constants of carbonic acid from potentiometric titrations. An iterative procedure described by Mehrbach et al. (1973) based on the Gran Method (Dyrssen and Sillen, 1967) was modified and used to calculate the equivalence points of the titration curve. K 1 * and K 2 * were then calculated as described by Goyet and Poisson (1989). During the potentiometric titration e.m.f. versus volume of acid added is recorded. Using the appropriate density of acid, the mass of acid added is calculated (W a ). The proton concentration [H + ] is calculated at each recorded e.m.f. using the Nernst equation. The weight of sample (m o ) and the hydrochloric acid concentration (C in mol kg -1 of solution) are known values. For the first iteration, an approximate value of W 2 (mass of acid added at the second equivalence point) is calculated by extrapolating the line from equation (2.8) to zero: W a = a + b F 1 (2.8) The coefficients a and b are determined using a linear least square procedure, where F 1 is defined for each titration point around the second equivalence point from: F 1 = TA m o W a C = [H + ] (W o + W a ) (2.9) when F 1 = 0, the total alkalinity in m o grams of sample is equal to the equivalents of acid added W a C. The intercept of a plot of W a versus F 1 yields W 2 and the product W 2 C represents the number of equivalents of alkalinity initially present in the seawater sample (Mehrbach et al., 1973). In the consecutive iterations, the definition of F 1 includes the effect of side reactions due to the presence of HCO - - 3, HSO 4 and HF in the solution: F1 = (m o + W a ) [H + + HSO HF - HCO - 3 ] (2.10)

42 23 in equation (2.10), HSO - 4, HF and HCO - 3 were calculated using: [HSO - 4 ] = [SO 2-4 ] Total (1 + K* (HSO4-) / [H + ]) -1 (2.11) [HF] = [F - ] Total (1 + K* (HF) / [H + ]) 1 (2.12) [HCO - 3 ] = TCO 2 (1 + [H + ] / K* 1 ) -1 (2.13) The value of TCO 2 used in equation (2.13) was calculated from: TCO 2 = C (W 2 W 1 ) (2.14) after the first approximation of W 1. After calculating W 2 in the first iteration, the value of W 1 (mass of acid added at the first equivalence point) is calculated by extrapolating W a = a + b F 2 (2.15) to zero. This equation is the result of a linear least square fit (a and b are respectively, the intercept and the slope), where F 2 is defined for each titration point around the first equivalence point (Mehrbach et al., 1973) by: F 2 = K 1 * (W a -W 1 ) = [H + ] (W 2 - W a ) (2.16) In the consecutive iterations the definition of F 2 included a correction due to the presence of carbonate ion in the solution (Mehrbach et al., 1973): F2 = [H + ] (W 2 - W a ) K1*/C {(W 2 - W a ) K 2 * C (1+2[H + ]/K 1 *) / ( (m o + W a ) [H + ] ) }(m o + W a ) (2.17) The first guess of K2* was computed from the regression of Mehrbach et al. (1973). In the consecutive iterations, K 2 * was computed as the average of the values from those points in the titration greater than zero, but lower than W 1 using the following equation (Goyet and Poisson, 1989):

43 24 K 2 * = [H + ] [{(W 1 - W a ) C [OH - ] (m o + W a ) } / { (W2 2W 1 + W a ) C + [OH - ] (m o + W a ) } (2.18) The concentration of [OH - ] can be calculated using the self-ionization of water (K w ): [OH - ] = K w* / [H + ] ) (2.19) The first dissociation constant of carbonic acid (K1*) was computed as the average of the values from those points in the titration between W 1 and W 2 using the following equation (Goyet and Poisson, 1989): K1* = [H + ] (W 2 - W a ) / (W a -W 1 ) (2.20) The first approximations of W1, W 2, K 1 *, K 2 * and TCO 2 were substituted into the more refined equations of F 1 and F 2 ; equations (2.10) and (2.21), respectively. The iterations were continued until successive values of K 1 * and K 2 * did not differ by more than 0.03%. The Gran method was also used to evaluate pk1* in natural seawater. In this case, the corrections for borate are included in the refined equation for F 2. [H + ]) F2 = [H + ](W 2 - W a ) K 1 * / C{(W 2 - W a ) K 2 * C(1+2[H + ]/K 1 *) / ((m o + W a ) + TB (1+[H + ]/K 1 *) / (1+[H + ]/K B ) } (m o + W a ) (2.21) where TB is the total borate and K B * is the dissociation constant of boric acid (K B * = [H + ][B(OH) 4 - ]/[B(OH)3 ]). C W i is the total number of chemical equivalents of the

44 25 weight species involved. The value of the second dissociation of carbonic acid K 2 * used in the former equation was taken from the regression of Mehrbach et al. (1973). There are several differences between the iterative procedure used for this work and the one used by Mehrbach et al. (1973): The standard potential E * used to start the iterations was 500 mv. The value of E * was refined during the first iteration and the later value was used in the rest of the iterations to calculate the concentration of [H + ] from the Nernst equation assuming a theoretical behavior of the electrode system. Mehrbach et al. (1973) calibrated their electrode systems using dilute buffers from the National Bureau of Standards and determined the apparent free hydrogen ion activity coefficient from the titrations at a given temperature and salinity. The corrections for borate were discarded in the refined equation for F modified Gran method was used to analyze titrations in artificial seawater. The second dissociation constant of carbonic acid K2* was included in the iteration procedure when the modified Gran method was used to analyze titrations in artificial seawater. Mehrbach et al., (1973) used (Lymann, 1956) values of K 2 * and K B * for the iteration procedure, which are not as reliable as recent estimates of K B * (Dickson, 1990a) and pk 2 * (Millero et al., 2006). 2, when the

45 The determination of K 1 *K 2 * and K 2 *. The pk 1 *+pk 2 * value can be determined from a TA titration on a closed cell by determining the product K 1 * K 2 * which is defined by K 1 *K 2 * = [H + ] 2 [CO 3 2- ] / [CO 2 ] (2.22) The carbonate alkalinity (CA) and total carbonate (TCO 2 ) are defined by CA = [HCO 3 - ] + 2[CO 3 2- ] (2.23) TCO 2 = [CO 2 ] + [HCO 3 - ] + [CO 3 2- ] (2.24) At the equilibrium ph for a solution with added NaHCO 3, CA is equal to TCO 2 (Mehrbach et al., 1973). This gives the following equation: [CO 2 ] = [CO 2-3 ] (2.25) and the product K 1 *K 2 * can be determined from K 1 *K 2 * = [H + ] 2 (2.26) The equilibrium ph 0 is given by ph 0 = ½ (pk 1 * + pk 2 *) (2.27) The ratio TCO 2 /CA is given (Mehrbach et al., 1973) by TCO 2 /CA = ([H + ] 2 + [H + ]K 1 * + K 1 *K 2 *) / ([H + ]K 1 * + 2K 1 *K 2 *) (2.28) The purity of the bicarbonate used in the experiments is TCO 2 /CA (pure bicarbonate has a value of one, A = 1). An error of ±0.002 in ph leads to an error of ± in TCO 2 /CA. This is much smaller than an error of in TCO 2 /CA due to an error of 0.01 in pk 1 *, as well as an error of in TCO 2 /CA due to an error of 0.02 in pk 2 *. This leads to an overall probable error of in TCO 2 /CA, which is higher than the experimental uncertainty in the ratio TCO 2 /CA (± ) for NaHCO 3

46 27 determined from direct measurements of TCO 2 and CA (Mojica-Prieto and Millero, 2002). The value of K 1 *K 2 * can be calculated from (Mehrbach et al., 1973): K 1 *K 2 * = (10-2 pho + [1 TCO 2 /CA ] K 1 *10 -pho ) / (2 TCO 2 /CA 1) (2.29) The ratio TCO 2 /CA for the sodium bicarbonate is critical for the determination of K 1 *K 2 *. An error of ± in TCO 2 /CA can lead to errors of ±0.007 in ½(pK 1 * + pk 2 *) and this would propagate to errors of ±0.015 in pk 2 *. The errors in ½(pK 1 * + pk 2 *) due to errors of in ph (0.001) and 0.01 in pk 1 * (0.001) are much smaller than those due to the errors in TCO 2 /CA (Mojica-Prieto and Millero, 2002). A typical determination of ½(pK 1 * + pk 2 *) using small additions of pure solid NaHCO 3 is given in Fig e.m.f. (mv) /2(pK 1 * + pk 2 *) Time (min.) Figure 2.1. Potentiometric determination of ½(pK 1 * + pk 2 *) using pure solid NaHCO 3, this is the steady state ph or ph 0.

47 28 The procedure to evaluate the product K 1 *K 2 * from the ph 0 measurements is the following. From the definition of the two dissociation constants of carbonic acid, the product of the dissociation constants can be written as in equation (2.22). After an entire titration with HCl and when all CO 2 was removed from a given sample, a small amount of NaOH was added to neutralize the solution. Then small amounts of NaHCO 3 (s) were added until the equilibrium conditions and the steady state ph 0 was reached. For more details see Mojica-Prieto and Millero (2002). At the steady state, the carbonate alkalinity and the total inorganic carbon are the same, equation (2.25) (Mehrbach et al., 1973). The product K 1 *K 2 * can be calculated as equation (2.26) and rewritten as equation (2.27). The ratio TCO 2 /CA, as given by Mehrbach et al. (1973), represents the purity of the bicarbonate used in the experiments. If the sodium bicarbonate is contaminated with carbonate, the value of TCO 2 /CA will be A < 1. This is usually the case, even with high purity bicarbonate. When the K 1 *, TCO 2 /CA and ph 0 values are known, the product of the two dissociation constants of carbonic acid K 1 *K 2 * can be calculated using equation (2.29). The ratio TCO 2 /CA from the bicarbonate is critical for the determination of K 1 *K 2 *. The second dissociation constant of carbonic acid K 2 * can be calculated by dividing the product K 1 *K 2 * by an appropriate value of the first dissociation constant K 1 * Fugacity of CO 2 (fco 2 ). The partial pressure of CO 2 was determined in seawater and in the atmosphere using a flowing system similar to the one designed by Wanninkhof and Thoning (1993). The equilibrator used during the Bahamas cruises was based on the design by Weiss

48 29 (1981) in order to make continuous measurements and modified by Goyet and Peltzer (1993). The fraction of CO 2 in the equilibrated air sample is measured using a differential, non-dispersive, infrared LI-COR LI-6262 CO 2 /H 2 O analyzer (Goyet and Peltzer, 1993). Samples were measured wet and the signal was corrected to dry air at 1 atm pressure using the measurements from the LI-COR water channel and connected pressure gauge. The water channel was calibrated using the LI-COR Dew Point Generator, model LI-610, as described in the LI-6262 Instrument operation manual. The reference chamber of the IR analyzer was continuously purged with dry, CO 2 -free air. This allowed the operation of the instrument in the absolute mode and gathered CO 2 concentrations directly from it. The system was calibrated using three standard gases 250, 340 and 500 ppm CO 2 in air. The net air-sea flux of CO 2 or pco 2, which is the difference between the CO 2 partial pressure in the atmosphere, and that in the surface ocean, is more precisely expressed as CO 2 fugacity (fco 2 ). fco 2 in a seawater sample is determined from the measured mole fraction of CO 2 gas corrected to dry air and to the pressure of 1 atm (X CO2 ) according to: fco 2 = [X CO2 P(1-p H2O /P)] exp[(b δ 12 )P/RT] (2.30) where P is the atmospheric pressure (atm); p H2O is the water vapor pressure at the temperature of the water in the equilibrator (atm), and T is the temperature of the sample in the equilibrator (K). The exponential term is the fugacity correction where B 11 is the second virial coefficient of pure CO 2 (B 11 = T T 3 ) and δ 12 is the correction for an air-sea mixture according to δ 12 = T (Weiss, 1974).

49 Discussion of the pk 1 * + pk 2 * method. The pk 1 * + pk 2 * method was an excellent procedure to analyze and quantify the carbonic acid pk 1 * and pk 2 *, with and without the addition of boric acid. To elucidate their ionic interactions, Pitzer equations were used to examine how different media affect the constants. The computer software (CO2Brine) gives an estimate of the values pk 1 * and pk 2 * values at different ionic strengths and temperatures with an input of two CO 2 parameters. The earlier pk 1 * measurements in NaCl solutions were made by Harned and Bonner, (1945). Measurements of pk 1 * and pk 2 * were made by Millero and Thurmond (1983) and He and Morse (1993). More recent measurements over a wide range of temperature and ionic strength were made by Millero et al. (2007). Measurements of the constants in ASW were made by Hansson (1973); Goyet and Poisson (1989); Roy et al., (1993a); Mojica-Prieto and Millero (2002). A better understanding of the interactions between B(OH) 3 and other electrolytes is needed. Chanson and Millero (2006) examined the interaction of B(OH) 3 with a number of electrolytes including some sea salts. The value of pk B* in seawater is close to the carbonic acid pk 2 * value, thus one might expect B(OH) 3 to interact more with CO 2-3 than HCO - 3. Many studies reported reliable values of pk 1 * in ASW with boric acid, but only a few were able to get reliable values that agreed with the values of pk 2 * in seawater (Mojica-Prieto and Millero, 2002). The differences in the pk 2 * values in real seawater and ASW have been attributed to organic material (potential unrecognized protolyte) or interactions with boric acid (Bradshaw and Brewer, 1988; Millero et al., 2002).

50 31 Computer software (TFM in Excel by Dr. Pierrot, RSMAS, now at CIMAS) fit the titrations of seawater based on earlier equations (Dyrssen and Sillèn 1967; Park, 1969; Dickson, 1981). When using this program there is difficulty in fitting ASW with and without boric acid. The TFM computer code was modified to make calculations in ASW with added boric acid. I also used the CO2Brine program (BASIC programming) to calculate the chemical equilibria. Newer programs are now available for different platforms to calculate and fit data, MATLAB (van Heuven et al., 2009) and for ocean acidification in R platform (Hofmann et al., 2010).

51 Experimental methods and equipment for the pk 1 * and pk 2 * from titrations with hydrochloric acid. Potentiometric titrations on a series of experiments in NaCl solution media at different ionic strengths from 0.25 to 2.0m, and changes in Na 2 CO 3 and B(OH) 3 from 250 to 2000μm were made to determine values of pk 1 * and pk 2 *. These measurements were made to examine how changes in the composition affect the carbonate constants. The effect of change in temperature and ionic strength were examined (t = 0-45ºC; I = M). These results are given in more detail in section Measurements of pk 1 * + pk 2 * from samples stripped of CO 2 and then with solid NaHCO 3 additions to a constant ph <steady state ph>. The ½(pK 1 * + pk 2 *) method (Mojica-Prieto and Millero, 2002) was used in order to get a more reliable estimation of pk 2 * values for oceanic CO 2 parameter determination. The ½ (pk 1 * + pk 2 *) method was applied to the following set of media: 1. Real seawater S = 36; t = 25 C 2. Sodium chloride solutions 0.7 M; t = 25 C 3. Regular formulation of artificial seawater S = 35, t = 25 C 4. Sodium chloride solutions 0.7 M; boric acid 400 μm; t = 25 C 5. Artificial seawater S = 35; boric acid 400 μm; different amounts of calcium and magnesium; t = 25 C 6. ASW at 25 C with up to 8 μm of organic material (with a pk* similar to the carbonic acid pk 2 *). 7. Tests 1-6 over a wide range of temperature (0-50 C)

52 Equipment and materials. The salinity of SW was measured with a Guildline 8410 PortaSal Salinometer. This instrument is regularly calibrated with standard seawater of known conductivity on the Practical Salinity Scale. A vibrating densimeter Mettler/PAAR DMA 60 was used to measure the density of the samples to a precision of kg m -3. All the solutions were prepared in molal concentrations (m) by weight, using ion exchanged (18 M Ω) water from a Millipore Super-Q system (milliq water). The balances used were the Mettler PM1200 and H20 models. The measurements were reliable to g and g, respectively. The titrations were made with automated titration systems (Brinkman Dosimat 645 and Orion ph-meter 720A) controlled by a PC. The solutions were titrated with 0.25 M HCl that contained 0.45 M NaCl. The acid batches used in this study #9601 and STD 03, had concentrations of M and M, respectively. The titration cells used were #16, #17 and #30. They were similar to those developed by Bradshaw and Brewer (1988) with improvements made by Millero et al., (1993). They were made of Plexiglas having volumes near 200 ml with a 5mL piston and openings for the 8101 Orion Ross glass ph electrode and the Orion double junction Ag AgCl reference electrode (Fig. 2.2). The systems were calibrated using certified reference materials (CRM) prepared by Dickson et al. (2003). During the experiments the e.m.f. was monitored with a ph meter and the temperature was monitored with a Guildline 9540 digital platinum resistance thermometer. The thermometer was calibrated on the NIST 1990 temperature scale and

53 34 controlled with a Neslab RTE 221 circulating bath to ±0.01ºC. The temperature of the HCl and sample was also controlled to a fixed temperature. The closed cell has a fill and a drain valve that allows a good reproducibility of the volume. An MS-Windows CVI program controlled the titration by recording the e.m.f. reading from the ph meter after the signal became stable (±0.05 mv). The titration ranged from 25 minutes at temperatures higher than 25 C to about 85 minutes at temperatures lower than 25 C. The computer program analyzes the data collected during the titration and calculates the total alkalinity (TA), the total dissolved inorganic carbon (TCO 2 ), the standard potential of the cell (E 0 ), the ph and the pk 1 *.

54 35 Dosimeter Burete ph meter Computer Reference electrode Glass Piston Overfill Valve Water jacket Glass electrode Acid delivery tip Magnet Drain-fill valve Water bath Magnetic plate Figure 2.2. Closed cell for the titration of NaCl solutions, ASW and SW with HCl ~0.25M. The performance of the titration system was monitored by comparison with CRM batches #57 and #61. At least three replicates were measured for each type of solution potentiometric titration. Typical titrations in different media are shown in Figs. 2.3 and 2.4.

55 e.m.f. (mv) Vol. HCl 0.25 M (ml) Figure 2.3. Typical titration for Total Alkalinity using HCl ( 0.25 N). A high resolution curve was setup for detail carbonic acid pk i * fittings in all NaCl solutions. For some of the titrations the BSTAC4 program which uses a Quick Basic platform was used to fit the titrations data and determine the carbonic acid pk 1 * and pk 2 * in molar units (DeStefano et al., 1993). All the dissociations constants were converted to the seawater scale (mol kg -1 soln).

56 Titration of ASW and real SW 400 e.m.f. (mv) Seawater Artificial Seawater Volume of HCl ( N) in ml Figure 2.4. Typical titration for Total Alkalinity in SW and ASW with Na 2 CO 3 added. A detailed high resolution curve is obtained from the titration setup to better fit the carbonic acid pk i * Reagents and solutions. A Na 2 CO 3 (s) reagent grade from Aldrich Chemical Co., A.C.S. reagent (99.999%) was used to avoid some undesired impurities that can cause problems; NaCl(s) A.C.S. reagent (99+%) from Aldrich Chemical Co. and Boric Acid(s) from J.T. Baker Chemical Co., A.C.S. reagent grade (99.6%) with impurities of Ca, Cl, SO 4 and PO 4 less than 0.004% each. The seawater used in this study is Gulf Stream seawater collected off the coast of Miami. The seawater was passed through a 122 mm in diameter and 0.45μm pore size filter made by PALL Supor -450 filters before it is used. The ASW without NaHCO 3 was prepared as described by Millero, (1996) according to the specifications of Goyet and Poisson, (1989).

57 The effect of organic carbon on the carbonic acid dissociation constants. Organic carbon content in artificial seawater has little or no effect on the carbonic acid pk 1 * (Roy et al., 1993b). The dissolved organic carbon (DOC) in seawater was changed to examine its affect on the carbonic acid dissociation constants, pk i *. As stated earlier, DOC could be the cause of the differences in pk 2 * in real and ASW (Bradshaw and Brewer, 1988; Millero et al., 2002). There have been some attempts to examine how organic constituents from the marine phytoplankton can affect TA (Fraga and Alvarez- Salgado, 2005; Hernández-Ayón et al., 2007; Kim and Lee, 2009). The effects are believed to be small but further study is needed Manipulation of seawater to gather information on carbonic acid dissociation constants Mixing of ASW with real seawater The mixtures of real seawater with ASW with and without boron allow one to examine the effect of boric acid and DOC on the carbonic acid constants. The concentration of boron in Gulf stream seawater at S = is 430µM (Uppström, 1974) and the new non conservative B/Cl ratio has a value of 446 µm (Lee et al., 2010) Addition of boric acid to real seawater For higher values of B T one can add increasing amounts of boric acid to real seawater. This gives information of how boric acid affect the values of pk 1 * and pk 2 * Reduction of the DOC load from real seawater An attempt was made to examine the effect of DOC on the pk 2 * values. This was done by using adsorption columns and irradiating seawater with a UV lamp. The SW

58 39 was passed through a Whatman GF/C (1.2µm) filter and then through a pair of C18 columns used to trap organic molecules from Dr. Zica s Laboratory (RSMAS). The SW was also irradiated for 12 hr with a UV radiation lamp of Hg. The DOC content was then measured with a High Temperature Catalytic Oxidation TOC (Shimadzu TOC-V CSH ) by Charles Farmer from the Organic Biogeochemistry Research Group (RSMAS-University of Miami). The DOC was also measured in Milli-Q water (MQW), the NaCl solutions, ASW and SW diluted with MQW.

59 Methods used in different media for carbonic acid pk i *. The artificial seawater (ASW) was prepared with Milli-Q water using high purity salts close to a salinity of 35 as described by Millero (2006) according to the proportions and methods of Goyet and Poisson (1989). This ASW solution resembles the average chemical composition of seawater better than the composition used by earlier authors (see Table 2.1). The gravimetric salts (NaCl, Na 2 SO 4, KCl, Na 2 CO 3, NaF and KBr, all reagent grades) were dried overnight in an oven at 60 C before weighing and addition to Milli-Q water. The hydrated salts (MgCl 2, CaCl 2 and SrCl 2 ) were added from stock solutions of known concentration. The concentrations of the stock solutions were determined from density measurements made with a Mettler/Paar Vibrating Density Meter DMA/60 and using the equations of state for the salts (Lo-Surdo et al., 1982). The density of the ASW solution at 25 C was kg m -3, which is in reasonable agreement with real seawater ( kg m -3 ) at 25 C and S = 35 (Millero and Poisson, 1981). The artificial seawater was stored at room temperature ( 25 C) in 20L HDPE or PP Nalgene Bottles before use. All the salinities were determined on a Guildline 8410 PortaSal Salinometer and using a DMA 60 Mettler/Paar Density Meter. The salinities were also determined from the density measurements using the 1 atm equation of state of Millero and Poisson (1981). The ASW TA and TCO 2 were determined with HCl using an automated potentiometric titration system (Millero et al., 1993a). The performance of the system was monitored using certified reference material (CRM) and locally prepared reference materials (RM). The reference materials indicated that the precisions were ± 2 µmol/kg for TA and ± 3 µmol/kg for TCO 2. The external reference was CRM Batch # 61 provided

60 41 by Dr. Andrew Dickson (Scripps Institution of Oceanography, La Jolla, CA). The internal references were RM Batches 5 and 6. These are batches from the Gulf Stream seawater collected off the coast of Miami. Table 2.1. Composition of artificial seawater in moles kg -1 solution used by various authors. Hansson Roy et al. Goyet & Poisson Seawater a NaCl Na 2 SO KCl CaCl MgCl NaF KBr SrCl Na 2 CO NaHCO B(OH) a ) (Millero, 1996; Millero, 2006; Millero et al., 2008). The composition of the ASW was identical to the formula used by Goyet and Poisson (1989). The main difference is that in some of the experiments we added boric acid and/or modified the fluorine content. All the samples were equilibrated to the desired temperature in the titration vessel using a Neslab RTE-221 water bath that controlled the temperature to ± 0.05 C. Flowing water at the desired temperature was circulated through the titration cell external chamber and around the piston delivering HCl during the measurements. The temperatures in the constant temperature bath and in the cell were measured with a Guildline 9540 Digital Resistance Thermometer. The temperature inside the cell was measured before and after each titration. The values agreed to ± 0.1 C which is

61 42 equivalent to an error of ± in pk 1 * and ± in pk 2 *. The recorded temperature of a run is the mean of the initial and final values. The titration system (Millero et al., 1993a) consisted of a closed water-jacketed plexiglass cell with a ROSS 8101 glass ph electrode and an Orion double junction Ag/AgCl reference electrode. Some measurements were made using an Orion 8005 ROSS Reference Half-Cell electrode. The titrant was delivered with a Metrohm 665 Dosimat titrator and the e.m.f. measured with an Orion 720A ph meter. The system was controlled by a personal computer (Millero et al., 1993a) using a National Instrument s Labwindows/CVI environment. The titration was made by adding ~0.25M HCl (in 0.45 M NaCl) to seawater past the carbonic acid end point. A typical titration records the e.m.f. after the readings become stable (± 0.05 mv) and adds enough acid to change the voltage to a pre-assigned increment (9 mv). This provides more data points in the range of a rapid increase in the e.m.f. near the endpoints. The values of pk 1 * and pk 2 * were determined for ASW using the methods of Mehrbach et al., (1973) and a non-linear curve-fitting procedure (Dickson, 1981; Johansson and Wedborg, 1982). This procedure was modified to a more user-friendly MS-Excel version by Dr. Denis Pierrot. The program determines the E*, pk 1 *, pk 2 *, TA and TCO 2 of the sample from the full titration (>50 pts). The dissociation constants needed in the computer code were taken from the literature: B(OH) 3 from Dickson (1990a); HF, from Dickson and Riley (1979); HSO - 4, from Dickson (1990b); H 2 O from Millero (1995). The calculations were carried out in a manner similar to the methods described by Goyet and Poisson, (1989). Care is needed in fitting titration data with many variables. To test the reliability of the computer software, a number of titrations as

62 43 a function of temperature were made on certified reference material (National-Research- Council, 2002; Dickson et al., 2003; Dickson, 2004a) provided by Dr. Andrew G. Dickson with known values of TA and TCO 2. The values of pk 1 * + pk 2 * were determined in ASW stripped of CO 2 after the addition of acid (HCl). The ph was then adjusted with 0.1 N NaOH within ± 0.05 ph of the equilibrium value ph = ½(pK 1 + pk 2 ). Solid NaHCO 3 was then added to the solution until a constant ph was reached. Unlike the study of Mehrbach et al. (1973), the electrode was calibrated in the sample by titrating it with HCl before the bicarbonate addition. The ph was also determined using the spectrophotometric technique with the indicator m-cresol purple (Clayton and Byrne, 1993). The values of pk 1 * were determined from potentiometric titrations as described below. The HCl (in 0.45 M NaCl) used for the potentiometric titrations was prepared in a large batch and stored in 500 cm 3 borosilicate glass bottles. The concentration of the HCl was determined by coulometry (Taylor and Smith, 1959). The batch used in this study (#9601) had a concentration of ± M. The NaHCO 3 used for the determination of the product K 1 *K 2 * was reagent grade of highest purity (Sigma- Aldrich A.C.S. Sigma Ultra, 99.5%, Batch #084K0190). The purity of the NaHCO 3 can be represented as A by the ratio of TCO 2 /CA. When the bicarbonate is the only chemical - species responsible for the alkalinity of the sample, the carbonate alkalinity (CA = HCO CO 2-3 ) and the total alkalinity (TA) of the solution are the same. Mehrbach et al. (1973) determined A from a comparison with a standard of KHCO 3. In this study, the ratio TCO 2 /CA was experimentally determined by independent measurements of TA by potentiometric titration and TCO 2 by coulometry. These experiments were performed in

63 44 a NaHCO 3 solution (0.002 M) in NaCl media (0.7 M) prepared with Milli-Q water devoid of CO 2. The measurements resulted in CA = ± 1.8 µmol/kg and TCO 2 = ± 0.8 µmol/kg. This yields a value of the ratio TCO 2 /CA = ± The value of CA has been corrected for the blank of 10 ± 1 µmol kg -1 determined in measurements of TA in 0.7 M NaCl with different amounts of NaHCO 3. The parameters involved in the calculation of TCO 2 /CA are given in the equation. TCO 2 /CA = ( [H + ] 2 + [H + ] K 1 * + K 1 * K 2 * ) / ( [H + ] K 1 * + 2 K 1 * K 2 * ) (2.31) where K 1 * and K 2 * are the values of the carbonic acid dissociation constants in NaCl 0.7 m solutions (pk 1 * = and pk 2 * = 9.516) from Millero et al., (2007). The values of TCO 2 /CA were used along with the measured TA to determine the contribution of the blank to the values of CA and TCO 2 in the bicarbonate, and were utilized to calculate the errors due to the blank. The 0.1 M NaOH solution used to adjust the ph before the bicarbonate additions was prepared from J.T. Baker s Dilute-It analytical concentrate diluted with Milli-Q water that was previously purged with a CO 2 free gas. The NaOH solution was kept free of CO 2 by using an Ascarite II trap (A.H. Thomas ). Boric acid B(OH) 3 was used in these studies instead of borax (Na 2 B 4 O 7 10H 2 O) as used in some of the previous ASW formulations. When borax is dissolved in water it forms the following chemical species according to: Na 2 B 4 O 7 10H 2 O 2B(OH) 3 + 2Na B(OH) 4 + 3H2 O (2.32) This yields an equimolal mixture of boric acid plus borate which requires a slight hydrolysis correction.

64 Summary In summary, details on the experimental methods for the following carbonate parameters were presented TA, TCO 2, ph, fco 2 and K 1 * K 2 *. The analytical reproducibility of each parameter was 2 µm kg -1 ; 2 µm kg -1 ; 0.005; 2 µatm; and 0.005, respectively.

65 Chapter 3 DISSOCIATION CONSTANTS OF CARBONIC ACID IN SEAWATER AS A FUNCTION OF SALINITY AND TEMPERATURE To examine the thermodynamics of the carbonic acid system in seawater from measurements of ph, total alkalinity (TA), total carbon dioxide (TCO 2 ) and the partial pressure of carbon dioxide (pco 2 ) reliable constants are needed for the dissociation of carbonic acid. In this chapter, recent measurements of these constants published in Millero et al., (2006) are discussed. 3.1 Modeling the carbonate system in natural waters. The differences in the values of pk i * near 25 C between ASW and SW can be attributed to changes in the activity of the H 2 O and CO 2 or the activity coefficients of HCO - 3 and CO 2-3. For example, the increase in K 1 * and decrease in K 2 * can be attributed to a decrease in γ HCO3 and an increase in γ CO3 between ASW and SW. If these effects are due to the boric acid in real seawater (Mojica-Prieto and Millero, 2002), they can be attributed to interactions of HCO - 3 with B(OH) 3 or B(OH) and interactions of CO 3 with B(OH) 3 or B(OH) - 4. It is also possible that, the differences are related to an organic acid present in all seawater (Millero et al., 2002), which causes pk 2 * differences in SW and ASW. Further measurements are needed to pinpoint the cause of the differences. The measured stoichiometric constants (K 1 * and K 2 *) for carbonic acid in seawater are related to the thermodynamic constants. The activity coefficients of H +, HCO - 3, CO 2-3 and activities of CO 2 and H 2 O can be determined from ionic interaction 46

66 47 models (Pitzer, 1991). Millero and Roy (1997) have developed a carbonate model valid from t= 0 to 50 C and I = 0 to 6m. The model considers the ionic interactions in solutions of the major components of seawater and other natural waters (H-Na-K-Mg-Ca-Sr-F-Cl-Br-OH-HCO 3 -B(OH) 4 - HSO 4 -SO 4 -CO 3 -CO 2 -HF-B(OH) 3 -H 2 O). It has been used to predict the activity coefficients of the major and minor components of ions required to determine the dissociation constants of all the acids needed to examine the carbonate system in natural waters (H 2 CO 3, B(OH) 3, H 2 O, HF, HSO - 4, H 3 PO 4, H 2 S, NH + 4 ). The predicted dissociation constants for a number of acids in seawater have been shown to be in good agreement with experimental measurements (Millero and Roy, 1997). The model can be used to examine the effects of composition on the carbonate constants in seawater and to compare the measurements made by various studies. The compositions of artificial seawater used in the various studies are given in Chapter 2, Table 2.1. A calculation of the values of pk 1 * and pk 2 * using the compositions shown in Table 3.2 indicate that the differences in pk 1 * and pk 2 * are all within ± which is well within the experimental error of the measurements (see Table 3.1). Changes in the 2- values of SO 4 and F - show the largest effects on the values of pk i *, (an increase of in F - increases the pk s by and an increase of in SO 2-4 increases the pk s by 0.005). These effects cannot account for the increase in pk 1 * and decrease in pk 2 * between SW and ASW near 25 C. A comparison of the model calculations of pk 1 * and pk 2 * at S = 35 and from 0 to 45 C with the fitted measurements of Mehrbach et al., (1973), Goyet and Poisson (1989), Roy et al., (1993b) and Mojica-Prieto and Millero (2002) is shown in Fig The

67 48 model calculations of pk 1 * are in reasonable agreement with all of the measurements from 10 to 45 C. The model calculations of pk 2 * are in agreement with the measurements in seawater by Mehrbach et al., (1973) and Mojica-Prieto and Millero (2002). Large offsets occur in the values of pk 2 * made in ASW by Goyet and Poisson (1989) and Roy et al., (1993b). The model appears to be in error below 10 C probably due to the scarcity of measurements of pk 1 * and pk 2 * in NaCl solutions at low temperatures. In summary, the model agrees with the measurements in real seawater by Mehrbach et al., (1973) and Mojica-Prieto and Millero (2002). pk 1 (Meas - Calc) Model Comparisons Mehrbach et al. Goyet and Poisson Roy et al. Mojica-Prieto and Millero pk 2 (Meas - Calc) Temperature ( o C) Figure 3.1. A comparison of the values of pk 1 * and pk 2 * from 0 to 40 C and S = 35 determined by various workers with the model of Millero and Roy (1997).

68 49 The measurements by Mehrbach et al., (1973) and Mojica-Prieto and Millero (2002) on SW were not made in dilute solutions, thus they may not give reliable constants for estuarine waters. This is shown in Figure 3.2 by comparing the measurements of pk 1 * and pk 2 * of Mehrbach et al., (1973) with the model at 25 C. In this chapter, measurements of real seawater over a wide temperature (1 to 50 C) and salinity (5 to 50) range are shown. The seawater results of this study and the literature results have been fitted to equations that are valid for all marine waters over a wide range of salinity and temperature.

69 Experimental methods and equipment. The measurements were made on Gulf Stream seawater that was diluted or evaporated from salinity near 36. The seawater was filtered through a 0.45 µm filter and stored at room temperature in 50L P.P. Nalgene Bottles before use. The low salinity samples were made by diluting SW with pure Milli-Q water (18 M Ω) and the high salinity samples were obtained by slowly evaporating the SW. All the salinities below 42 were directly determined on a Guildline 8410 PortaSal Salinometer. The samples at higher salinities were determined by diluting the samples to salinities below 42 and measuring the density on a DMA 60 Mettler/Paar Density Meter. The salinities were determined from the density measurements using the 1 atm equation of state of Millero and Poisson (1981). For samples with salinity less than 8, sodium carbonate (0.002m) was added to aid in the determination of pk 2 *. Since the ph of seawater (~8) is not high enough to determine an accurate value of pk 2 *, small amounts of sodium hydroxide were added to increase the ph to ~9-10. These small additions of NaOH did not significantly change the salinity of the samples. The addition of NaOH was not necessary in dilute solutions when sodium carbonate was added.

70 51 pk 1 * pk 2 * Carbonic Acid Model Mehrbach et al S 0.5 Model Mehrbach et al. Figure 3.2. A comparison of the values of pk 1 * and pk 2 * at 25 C as a function of the square root of salinity (S = 0 to 45) of Mehrbach et al., (1973) with the model results from Millero and Roy (1997). All the samples were equilibrated to the desired temperature in a Neslab RTE-221 constant temperature water bath to ± 0.05 C before addition to the titration vessel.

71 52 Flowing water at the desired temperature was circulated through the titration cell and around the piston delivering the HCl during an experiment. The temperatures in the constant temperature bath and in the cell were measured with a Guildline 9540 Digital Platinum Resistance Thermometer calibrated on the NIST temperature scale by the manufacturer. The temperature inside the cell was measured before and after each titration. The values agreed to ± 0.1 C which is equivalent to an error of ± in pk 1 * and ± in pk 2 *. The recorded temperature of a run is the mean of the initial and final values. The titration system (Millero et al., 1993a) consists of a closed water jacketed plexiglas cell with a ROSS 8101 glass ph electrode and an Orion double junction Ag/AgCl reference electrode. Some measurements were made using an Orion ROSS Reference Half-Cell electrode. The ph electrodes were calibrated for each individual titration with the E* value. The titrant is delivered with a Metrohm 665 Dosimat titrator and the e.m.f. is measured with an Orion 720A ph meter. The system is controlled by a personal computer (Millero et al., 1993a) using a National Instrument s Labwindows/CVI environment. The titration is made by adding ~0.25M HCl (in 0.45 M NaCl) to seawater past the carbonic acid end point. A typical titration records the e.m.f. after the readings become stable (± 0.05 mv) and adds enough acid to change the voltage to a pre-assigned increment (9 mv). This provides more data points in the range of a rapid increase in the e.m.f. near the endpoints. More details on methods are given in Chapter 2, section The values of pk 1 * and pk 2 * [mol (kg soln) -1 ] were determined using a nonlinear curve-fitting procedure developed by Dickson (1981) and Johansson and Wedborg

72 53 (1982). This procedure was modified to a more user-friendly Excel version by Dr. Pierrot. The program determines the E*, pk 1 *, pk 2 *, TA and TCO 2 of the sample from the full titration (>50 pts). The electrodes are calibrated over the entire range of the titration with the computer code giving a value of E* that is constant at a given temperature and salinity. The dissociation or association constants needed in the computer code were taken from the literature. B(OH) 3 from Dickson (1990a); HF from Dickson and Riley (1979): HSO - 4 from Dickson (1990b) and H 2 O from Millero (1995). The dissociation constants on the seawater ph scale (Dickson, 1991) [H + ] SWS = [H + ] + [HSO - 4 ] + [HF] (3.1) where the brackets represent concentrations, mol (kg soln) -1 (Millero, 1995). The total seawater scale is given by [H + ] T = [H + ] + [HSO - 4 ] (3.2) The two ph scales are related by (Dickson, 1981; 1984) ph SWS = ph T + log ( 1 + β HSO4 [SO 4 ] T + β HF [F] T ) (3.3) where the subscript T represents the total concentration and βi are the association constants for the formation of HSO 4 - (Dickson, 1990b) and HF (Dickson and Riley, 1979) at the ionic strength of the mixture. The calculations were carried out in a manner similar to the methods described by Goyet and Poisson (1989). As pointed out by these authors, care is needed in fitting titration data with many variables. To test the reliability of the computer software, a number of titrations as a function of temperature were made on certified reference material provided by Dickson (2004b) with known values of TA and TCO 2. The derived values of pk 1 * and pk 2 * on these samples are given in Table 3.1. The values of pk 1 *

73 54 determined with and without the known values of TA and TCO 2 (floating) gave similar results that were in good agreement with the values of Mehrbach et al., (1973). The floating values of TA and TCO 2 are also in good agreement with the certified values. In a second series of studies, we examined the values of pk 2 * obtained with and without the addition of NaOH. These results indicate that to obtain values of pk 2 * that agree with the results of Mehrbach et al., (1973) or Mojica-Prieto and Millero (2002), an increase of the ph and the concentration of the CO 2-3 ion is needed. This was done in most of the experiments at high salinity. 3.3 Results and calculations The titrations were made on seawater samples as a function of temperature (1 to 50 C) and salinity (5 to 50). The measured values of pk 1 * and pk 2 * determined from the titrations are given in Table 3.2. The results are the average of duplicate or triplicate measurements on the same sample that agreed to within ± for pk 1 * and ± for pk 2 *. The values of pk 1 * and pk 2 * as a function of I 0.5 from 10 to 50 C are compared to the model calculations (Millero and Roy, 1997) in Fig The measurements are in good agreement with the model over this temperature range and approach the pure water values in dilute solution. As stated earlier, this is not the case at temperatures below 10 C at high ionic strengths.

74 55 Table 3.1. The effects of TA (µmol kg -1 ) and TCO 2 (µmol kg -1 ) levels and ph on the determination of the pk 1 * and pk 2 * at various temperatures. A. Effect of fixing or floating TA and TCO 2 on the calculated pk 1 * and pk 2 * Fixed Values of TA and TCO 2 Floating Values of TA and TCO 2 Temp Salinity TA TCO 2 pk 1 * pk 2 * ΔpK 1 * ΔpK 2 * ΔTA ΔTCO 2 20 o C Floating Values of TA and TCO 2 20 o C B. Effect of initial ph on the values of pk 1 * and pk 2 * Without NaOH With NaOH From Mehrbach Data Δ pk 1 * Δ pk 2 * Temp Salinity pk 1 * pk 2 * pk 1 * pk 2 * Without With a a) Certified Reference Material.

75 Table 3.2a. Measured values of pk1* and pk2* for carbonic acid in seawater as a function of salinity and temperature. Temp S pk1* pk2* Temp S pk1* pk2* Temp S pk1* pk2*

76 Table 3.2b. Measured values of pk1* and pk2* for carbonic acid in seawater as a function of salinity and temperature. Temp S pk1* pk2* Temp S pk1* pk2* Temp S pk1* pk2*

77 Table 3.2c. Measured values of pk1* and pk2* for carbonic acid in seawater as a function of salinity and temperature. Temp S pk1* pk2* Temp S pk1* pk2* Temp S pk1* pk2*

78 59 Table 3.2d. Measured values of pk1* and pk2* for carbonic acid in seawater as a function of salinity and temperature. Temp S pk1* pk2* Temp S pk1* pk2* Temp S pk1* pk2* a) The blank entries in the pk2* column were experiments where NaOH was not added and hence pk2* could not determined reliably.

79 pk 1 * Model 10 o C 20 o C 30 o C 40 o C 50 o C pk 2 * Figure 3.3. The measured values of pk 1 * and pk 2 * from 10 to 50 C as a function of the square root of ionic strength (I). The smooth curves are the values calculated from the model of Millero and Roy (1997). I 0.5

80 61 To examine the internal consistency of the measurements made at each temperature the results were first fitted to equations of the form pk i * pki 0 = A S B S + C S 2 (3.4) pk i * pki 0 = A I B I + C I 2 (3.5) where the ionic strength I = S / ( S), and the pk 0 i values are determined from Harned and Bonner (1945), and Harned and Davis (1943) as refit by Millero (1979) pk 0 1 = /T ln T (3.6) pk 0 2 = /T ln T (3.7) The average standard deviations for the individual temperatures varied from to for pk 1 * (weighted average of , N = 466) and from to for pk 2 * (weighted average of , N =458). All of the measurements as a function of temperature and salinity have been fitted to equations of the form: pk * i - pk 0 i = A i + B i / T + C i ln T (3.8) The adjustable parameters have been fitted to functions of salinity using equations: A i = a 0 S a 1 S + a 2 S 2 (3.9) B i = a 3 S a 4 S (3.10) C i = a 5 S 0.5 (3.11) Similar equations as a function of ionic strength can be formulated by replacing S with I.

81 62 The coefficients used were derived by using an F-test and are shown in Table 3.3 along with the standard errors of the fits (σ = for pk 1 * and σ = for pk 2 *). The differences between the measured and calculated values of pk 1 * and pk 2 * are shown in 2σ deviations of and 0.020, respectively (Figure 3.4 and 3.5) σ = σ pk σ Salinity σ = σ pk σ Temperature ( o C) Figure 3.4. The differences between the measured and fitted values of pk 1 * as a function of salinity and temperature.

82 σ = σ pk σ Salinity σ = σ pk σ Temperature ( C) Figure 3.5. The differences between the measured and fitted values of pk 2 * as a function of salinity and temperature.

83 64 Table 3.3. Coefficients for the fits of the values of pk 1 * and pk 2 * in seawater as a function of temperature, salinity and ionic strength. pk 1 * Coeff. Salinity ( S ) Ionic Strength ( I ) pk 2 * Coeff. pk 1 * Coeff. pk 2 * Coeff. S 0.5 a S a S 2 a E S 0.5 /T a S/T a S 0.5 ln T a Std. Error Number Note: in the case of ionic strength, replace S for I in the first column. 3.4 Discussion. Comparisons of pk 1 * and pk 2 * calculated from equations (3.8) to (3.11) with earlier workers are shown in Figure 3.6 and 3.7. Our measurements for pk 1 * at S = 35 are in good agreement with those of Mehrbach et al., (1973), Roy et al., (1993b), Goyet and Poisson (1989), and Mojica-Prieto and Millero (2002) from 0 to 40 C. The measurements of Mojica-Prieto and Millero (2002) are higher than the other studies above 30 C. Our pk 2 * at S = 35 are in good agreement with those of Mehrbach et al., (1973), and Mojica-Prieto and Millero (2002) from 0 to 45 C. These comparisons demonstrate the reliability of our measurements.

84 65 pk 1 (Ours - Others) Mehrbach et al., 1973 Roy et al Goyet and Poisson, 1989 Mojica-Prieto and Millero, pk 2 (Ours - Others) Temperature ( C) Figure 3.6. A comparison of our pk 1 * and pk 2 * results at S = 35 with literature values.

85 66 pk 1 (Meas - Calc) Hansson, 1973 Mehrbach et al., 1973 Goyet and Poisson, 1989 Roy et al., 1993 This Study σ = Temperature ( C) 0.02 pk 1 (Meas - Calc) σ = Salinity Figure 3.7. Comparison of this work for pk 1 * and literature values as a function of temperature and salinity. The standard error of the fit for all the measurements is shown (see Table 3.4).

86 67 Since a number of the measurements on pk 1 * and pk 2 * from the literature are in reasonable agreement with our results, we have decided to derive equations that include all the reliable studies. This includes the literature values of pk 1 * from Mehrbach et al., (1973), Goyet and Poisson (1989), Mojica-Prieto and Millero (2002), and the values of pk 2 * from Mehrbach et al., (1973) and Mojica-Prieto and Millero (2002). The results for the coefficients for equations (3.8) to (3.11) are tabulated in Table 3.4 along with the standard error of the overall fits. The differences between the measured and calculated values of pk 1 * and pk 2 * are shown in Figs 3.7 andfigure 3.8. Most of the deviations are within 2σ. The standard errors of the individual workers from the overall equation are given in Table 3.5. Table 3.4. Coefficients for the fits of the values of pk 1 * and pk 2 * in seawater as a function of temperature and salinity (measurements combined with literature studies). pk 1 * pk 2 * Parameter Coeff. Coeff. S 0.5 a S a S 2 a 2-4.0E E-04 S 0.5 /T a S 0.5 lnt a S/T a Std. Error Number

87 68 This Study pk 2 (Meas - Calc) Mehrbach et al Mojica-Prieto and Millero, Temperature ( o C) pk 2 (Meas - Calc) σ = Salinity Figure 3.8. The differences between all the measured and fitted values of pk 2 * as a function of temperature and salinity.

88 69 Table 3.5. Comparisons of the standard deviations of the authors and the standard error of the fit for all the measurements. Author σ (pk 1 *) No. σ (pk 2 *) No. This Study Mehrbach et al Goyet and Poisson Roy et al Mojica-Prieto and Millero All These combined equations represent the dissociation constant of carbonic acid over a wide range of salinity and temperature. They can be used to examine the thermodynamics of the carbonate system in most estuarine and marine waters. Our equations assume that seawater is diluted with pure water. This may not be the case for some estuarine systems. If the composition is known, one can use Pitzer models (Millero and Roy, 1997; Millero and Pierrot, 1998) to account for the difference in the composition of an individual river. Since the dissociation constants are frequently used to calculate the parameters (ph, TA, TCO 2 and fco 2 ) controlling the CO 2 system in natural waters, it is important to examine the errors involved in these calculations using various inputs. This was done by examining the calculations of two of the unknown parameters with an input of ph-ta, ph-tco 2, fco 2 -TA and fco 2 -TCO 2. The calculations were made at S = 35 and t = 25 C at ph = 8, TA = 2400 µmol kg -1 and two levels of fco 2 = 350 and 1400 µatm and TCO 2 = 2050 µmol kg -1. The results are tabulated in Table 3.6 for errors of in pk 1 * and in pk 2 *.

89 70 Table 3.6a. Errors in the determination of CO 2 parameters at S = 35 and t = 25 C due to errors of in pk 1 * a Uncertainties due to error in pk 1 * Input Variables fco 2 ΔTCO 2 ΔTA ΔpH ΔfCO 2 (µatm) (µmol kg -1 ) (µmol kg -1 ) (µatm) ph-ta ph-tco fco 2 -TA fco 2 -TCO TA-TCO Table 3.6b. Errors in the determination of CO 2 parameters at S = 35 and t = 25 C due to errors of in pk 2 *. a,b Uncertainties due to errors in pk 2 * Input Variables fco 2 ΔTCO 2 ΔTA ΔpH ΔfCO 2 (µatm) (µmol kg -1 ) (µmol kg -1 ) (µatm) ph-ta ph-tco fco 2 -TA fco 2 -TCO TA-TCO a) Initial input for fco 2 = 350 µatm (ph = 8.1, TA = 2400 µmol kg -1, TCO 2 = 2050 µmol kg -1 ) and for fco 2 = 1400 µatm (ph = 7.585, TA = 2400 µmol kg -1 and TCO 2 = 2305 µmol kg -1 ). b) The total probable error can be estimated from the square root of the sum of the errors due to pk 1 * and pk 2 *.

90 Summary. The difference between our equations of pk 1 * and pk 2 * with earlier workers and the new values for real and artificial seawater at S = 35 and temperatures from 0 to 40 o C are shown in Figure 3.9 and Our measurements for pk 1 * and pk 2 * are in good agreement with the earlier work of Mehrbach et al., (1973) and Mojica-Prieto and Millero (2002). The values of pk 1 * in seawater from 10 to 35 o C are higher by as much as 0.01 as compared with the values determined in artificial seawater. The values of pk 2 * in seawater from 10 to 35 C are higher by as much as Attempts to understand the causes of these differences are discussed in chapter pk 1 * (others - Millero et al., 2006) Mehrbach et al., SW Hansson ASW Goyet and Poisson ASW Roy et al., ASW Mojica-Prieto and Millero SW Temperature C Figure 3.9. Comparison of the new equation for pk 1 * (Millero et al., 2006) with earlier workers at Salinity = 35. Mehrbach et al., (1973); Hansson, (1973); Goyet and Poisson, (1989); Roy et al., (1993a); Mojica-Prieto and Millero, (2002).

91 (others - Millero et al., 2006) pk 2 * Hansson Mehrbach et al., Dickson and Millero Goyet and Poisson Roy et al., Temperature C Figure Comparison of the new equation for pk 2 * (Millero et al., 2006) with earlier workers at Salinity = 35. Artificial Seawater measurements of Hansson, (1973); Goyet and Poisson, (1989) and Roy et al., (1993a). Seawater measurements of Mehrbach et al., (1973) and other works by Dickson and Millero (1987).

92 Chapter 4 EFFECT OF BORIC ACID AND DISSOLVED ORGANIC CARBON ON THE DISSOCIATION OF CARBONIC ACID To understand the behavior of carbonic acid in natural waters accurate thermodynamic properties of the carbonate system are needed (Park, 1969; Millero, 1979; Brewer et al., 1997; Millero et al., 2002; Dickson, 2004b; Millero et al., 2006). Many workers have used various formulas to examine the thermodynamic properties of artificial seawater (ASW) (Kester et al., 1967; Bidwell and Spotte, 1985). In this chapter, I examine the effect of boric acid (HB) and dissolved organic carbon (DOC) on the values of pk 1 * and pk 2 * in real (SW) and ASW. The aim of this research is to examine how changes in HB and DOC affect the dissociation constants of carbonic acid using well established titrations (Millero, 1995; Millero, 2007) and ph(0) = ½(pK 1 *+pk 2 *) measurements (Mehrbach et al., 1973; Mojica-Prieto and Millero, 2002) in ASW. A summary of measurements of the stoichiometric constants for the dissociation of carbonic acid made in real seawater (Mehrbach et al., 1973; Millero et al., 2002; Mojica-Prieto and Millero, 2002; Millero et al., 2006) and artificial seawater (Goyet and Poisson, 1989; Roy et al., 1993; Mojica-Prieto and Millero, 2002) is given in Table 1.2. Comparisons of pk 1 * in ASW and SW by various workers are shown in Chapter 1, Figure 1.4. These comparisons are made relative to the measurements of Mehrbach et al., (1973) as reformulated by Dickson and Millero (1987). The pk 1 * measurements in SW by Mehrbach et al., (1973) and in ASW by Goyet and Poisson (1989) and Roy et al., (1993) are in good agreement (within 2σ = 0.014). The pk 2 * results of Hansson (1973) in ASW do not agree with the measurements of Goyet and Poisson (1989) and Roy et al., 73

93 74 (1993b). The pk 2 * results in SW do not agree with the results in ASW above 5 C (Fig. 1.4). These results suggest that measurements of pk 2 * in SW and ASW are different. Internal consistency studies (Millero et al., 1993b; Wanninkhof et al., 1999; Lee et al., 2000; Millero et al., 2002) using field data and laboratory measurements of seawater (Lee et al., 1996; Lueker et al., 2000; Mojica-Prieto and Millero, 2002) have shown that the Mehrbach et al., (1973) pk 2 * measurements in seawater are more reliable than the measurements made in ASW. The extensive pk 1 * + pk 2 * measurements of Mojica-Prieto and Millero (2002) and Millero et al., (2006) yield values of pk 2 * that are in excellent agreement with the results of Mehrbach et al., (1973). The difference in the pk i * constants in SW and ASW has been attributed to boric acid interactions with the carbonate ion (Mojica-Prieto and Millero, 2002) or to the presence of an organic protolyte in seawater (Bradshaw and Brewer, 1988; Millero et al., 2002). To unravel these possibilities, accurate measurements of the carbonic acid stoichiometric equilibrium constants are needed. The observed decrease in K 2 * in SW can be attributed to an increase in γ CO3 in SW. If these effects are due to the boric acid in real seawater, a possible interaction of CO 2-3 with B(OH) 3 or B(OH) - 4 may be present. The aim of this chapter is to investigate the causes of the differences in pk 2 *.

94 Boric acid effect on the dissociation constants of carbonic acid. The first set of experiments to examine the effect of boric acid on natural seawater were conducted on modified Gulf Stream seawater (S = ) by the addition of weight amounts of boric acid up to 4 times the normal concentration in seawater. The results are given in Table 4.1 and are shown as a function of added boric acid (Fig. 4.1). The pk 2 * values appear to fit a second degree function of added HB. Although there is scatter, the pk * 2 appear to extrapolate to the ASW results without boric acid. Also, high deviations of pk * 2 from the Millero et al., (2006) equation for SW appear at higher loads of HB (Fig. 4.2). These measurements were made using the ph(0) = ½(pK 1 *+pk 2 *) method described earlier (see Chapter 2, section 2.3.1). The ASW and SW values of pk 1 * were taken from the titration using the software TFM (described in Chapter 2). The average results in ASW without added HB (pk 2 * = 8.93) are in reasonable agreement with SW (pk 2 * = 8.96). The results show that HB affects the difference in pk 2 * between real and artificial seawater. Seawater and artificial seawater are multi-electrolyte solutions (Millero, 1974; Whitfield, 1975), with a considerable number of chemical species to be taken in account for their interactions.

95 pk* Boric acid (HB) µm Figure 4.1. Effect of boric acid (HB) on real seawater carbonic acid pk 2 *. The red dot is the expected pk 2 * (8.99) when HB = 0. The regression is pk 2 *= E -4 * HB E -7 * HB 2.

96 77 Table 4.1. Results of the boric acid additions to Gulf Stream seawater (S=36.101) carbonic acid dissociation constants at t = 25 C. TA TCO 2 pk µmol kg -1 µmol kg -1 1 * pk 2 * HB µm pk* HB (µm) Figure 4.2. The difference in the value of pk 2 * with added boric acid in seawater compared to the pk 2 * values when HB=0 from Figure 4.1.

97 Effect of DOC on the values of pk i * in seawater. In order to quantify the effect of DOC and boric acid, measurements of carbonic acid dissociations constants were made at 25 C on diluted Gulf Stream (DOC ~ 65µM) and Biscayne Bay (DOC ~175 µm) seawater using Milli-Q water. The dilution allows us to reduce the DOC level down to 25 µm (the lowest is 18 µm for pure ASW at S=36 and 2.3 µm for our Milli-Q water). Regular potentiometric titration plus the ½(pK 1 + pk 2 ) methods were performed on all runs. There was no clear evidence of the DOC effect at regular or diluted concentrations on the carbonic acid constants in real seawater within our experimental error (±0.005). The DOC results of pk 2 * (Measured Millero et al., 2006) from Gulf Stream diluted seawater in a DOC range of 25 µm to 105 µm (solid blue dots) show no trend and the data was within the experimental error (Fig. 4.3). A second attempt was made with 18 C columns to reduce DOC from Biscayne Bay waters. The SW sample was UV irradiated; reducing DOC by 30%. Further measurements are needed to examine the effect of high levels of DOC in estuarine waters to improve upon these preliminary results. The mixtures of Gulf Stream seawater and ASW allowed for the measurement of the dissociation constants of carbonic acid at different loads of DOC and HB (Table 4.2). The DOC effect on pk i * of mixtures of SW and ASW were within the experimental error and no trend was detected (n=24; σ = 0.003). Nevertheless, an extensive study of the carbonate constants useful to estuarine waters in the total and free ph scales is available (Millero, 2010). All DOC measurements were made on a Shimadzu TOC-V CSH by Charles Farmer at RSMAS - University of Miami.

98 Diluted seawater pk 2 * Gulf Stream Seawater Biscayne Bay Seawater DOC µm Figure 4.3. DOC effect on the carbonic acid pk 2 * value. Solid dots are diluted seawater from the Gulf Stream. pk 2 * Open circles are natural seawater from Gulf Stream or Biscayne Bay (Measured Millero et al., 2006). The comparison with other pk i* data (Mojica-Prieto and Millero, 2002) on ASW is included in Table 4.3.

99 80 Table 4.2. Dissociation constants of carbonic acid at 25 C in mixtures of Gulf Stream seawater and ASW with different loads of DOC and HB (TA and TCO 2 in µmol kg -1 ). Salinity constant at SW/ASW HB DOC mix % µm µm TA TCO 2 pk 1 * pk 2 * 50/ / / / / / / / / / / / / / /

100 Results of carbonic acid pk i * in NaCl solutions. Measurements were made at 25 C on NaCl solutions mainly 0.7 m and few on 0.25 m (on which the internal reference solution of the electrode system was modified) with carbon concentrations (C T ) of 1000 µm and 2000 µm added as Na 2 CO 3 and different HB concentrations up to 2000 µm. The fitted equation for the dependence of pk 1 * on HB (µm) on a 0.7 m NaCl solution is: pk 1 * = * HB (4.1) The effect of HB on NaCl solutions with Na 2 CO 3 is measurable since the concentration of HB is high enough to produce ionic interactions with the carbonates in the system. In Fig. 4.4, a trend line on pk 1 * is observed for a 0.7 m NaCl solution at 25 C with 1000µm Na 2 CO 3. For the same conditions with instead Na 2 CO 3 at 2000 µm, a linear trend is also observed (Fig. 4.5). Nevertheless, in NaCl 0.7 m solutions, there is no evident effect of boric acid on pk 2 * at C T = 1000 µm or C T = 2000 µm. However, a linear trend is shown (Fig. 4.6).

101 pk* 1 = E -3 * HB pk* HB µm Figure 4.4. Values of pk 1 * in sodium chloride 0.7 m solution with Na 2 CO µm at 25 C over different concentrations of boric acid. The estimated standard error of the pk 1 * value is and for HB = 0, the estimated pk 1 * value is pk 1 * = E-5 *HB pk 1 * HB µm Figure 4.5. Values of pk 1 * on sodium chloride 0.7 m solution with Na 2 CO µm at 25 C over different loads of boric acid. The standard error of for the pk 1 * estimate is very good, considering this large amount of HB.

102 pk 2 *= E -5 *HB pk 2 * HB µm Figure 4.6. Values of pk 2 * on sodium chloride 0.7 m solution with 2000 µm Na 2 CO 3 at 25 C over different loads of boric acid. The standard error of the estimate is for pk 2. The pk 2 * = when no boric acid is added. A new set of experiments was focused on the temperature and total boron effect on the carbonic acid dissociation constants in NaCl solutions of 0.25 m and 0.7 m. The temperature range was from 5 C 45 C. The boric acid concentration range was from 0 µm to 2000 µm. The Na 2 CO 3 concentration range was from 0 µm to 2000 µm. Figures 4.7 and 4.8 show the temperature effect on the carbonic acid pk i * at a constant HB.

103 pk* 1 = E -3 * t pk* Temp. C Figure 4.7. Temperature effect on carbonic acid pk* 1 in 0.25 m NaCl solution with added carbonate (250 µm) and HB (250 µm). The standard error of the estimate is and the pk 1 *(0) = pk 2 * = E -3 * t E -5 t pk 2 * Temp. C Figure 4.8. Temperature effect on carbonic acid dissociation pk 2 * in 0.25 m NaCl solution with 250 µm Na 2 CO 3 and 250 µm boric acid. The standard error of the estimate is The pk 2 *(0) estimate at T = 0 C is

104 Discussion No effect of DOC on pk 2 * was observed from the direct dilution and the DOC reduction treatment using the 18 C + UV lamp treatment. The effect of fluorine on carbonic acid pk i also should be tested in ASW. In this work, a few experiments were made without fluoride and a great number with 70 µm fluoride concentration. Millero et al., (2007) made measurements of the pk 1 * and pk 2 * in NaCl solutions from 0 to 250 C to obtain the Pitzer coefficients used to determine the activity coefficients of CO 2, NaHCO 3 and Na 2 CO 3 over a wide range of concentrations. These coefficients are useful to extend the ionic interaction model for the carbonate system in hydrothermal brines. A search to identify an organic base in seawater that affects seawater total alkalinity, was made by Bradshaw and Brewer, (1988). They propose a concentration of about 8 µm and a pk 4; this was tested by Takahashi on pco 2 (Millero et al., 2002) in the pk 2 *. A search for that specific organic compound with these characteristics in seawater is hard to find, and more with a condition in concentration of 8 µm when the level of dissolved organic carbon (DOC) in deep waters is about 50µM (Fasham et al., 2001). If this exists, it needs first to be identified, and then measure for the affect on the carbonic acid dissociation constants. The measurements need to be checked with the calculations of TA from ph and DIC using the CO2sys software. If the presence of boric acid affects the carbonic acid pk 1 * and pk 2 * values in seawater media, there also should be an interaction that affects the activity coefficients of bicarbonate, carbonate and borate too. These also will change the pk B * value.

105 Summary. Carbonic acid stoichiometric dissociation constants were determined in different media to test the effects of composition. The two areas of interest include dissolved organic carbon (DOC) and boric acid and how changes in their composition affect the pk * i. The boric acid effect on the carbonic acid pk * 1 in seawater were minimal. The * additions of small amounts of boric acid to ASW were in agreement with the earlier pk 2 measurements of Mojica-Prieto and Millero (2002). Further additions of boric acid to seawater, caused an increase in the pk 2 *. Similar measurements in NaCl showed decreases in both pk 1 * and pk 2 *. These were different than the results in seawater and are apparently related to the interactions of Mg 2+ and Ca 2+ in seawater. Further studies are needed to elucidate these effects. As discussed in earlier studies, the effect of boric acid on the value of pk 2 * in seawater was thought to be due to its affects on the activity coefficient of the CO 2-3 (Mojica-Prieto and Millero, 2002). Attempts to determine the effect of changes in DOC concentration were minimal and need further measurements. Studies on the effect of additions of humic acid of known DOC to seawater, may be useful in elucidating the effect of DOC on the values of pk * i.

106 87 Table 4.3. Diluted ASW results on steady state ph 0, ½(pK 1 +pk 2 ) and carbonic acid stoichiometric dissociation constants. Included the results of ph 0 for Mojica-Prieto and Millero (2002) = M&M. Temp. TA TCO Salinity 2 ph pk C µm kg -1 µm kg 1* pk 2* ph -1 0 ½(pK 0 1+pK 2) M&M

107 Chapter 5 THE FORMATION OF WHITINGS ON THE LITTLE BAHAMA BANK 5.1 Background. Whitings or whiting events (Bathurst, 1975) are suspensions of fined-grained carbonate minerals. Potential sources of the suspended carbonate minerals may be lowmagnesium calcite, high-magnesium calcite, and aragonite. These minerals include carbonates produced by physical and biological disintegration of animal and algal bioclasts, mineralizing pelagic organisms, blooms of microscopic algae, spontaneous formation of whitings from either abiotic precipitation or calcification of suspended picoplankton and organic matter. Whitings occur in lakes (Brunskill, 1969), and tropical and subtropical seas (Morse et al., 1984, 2003; Shinn et al., 1989; Robbins and Blackwelder, 1992; Millero et al., 2005). Whitings are relatively rare in the marine environment. The origin of the suspended calcium carbonate, which is dominantly aragonite with a minor amount of high magnesian calcite, found in marine whitings has been a source of controversy. Whitings on the Great Bahama Bank (GBB) are the most extensively investigated, having been studied since 1940 (Smith, 1940a, b). The source and cause of these whitings have been hotly debated for years (see Morse et al., 2003, for review, discussion and references). Morse et al., (2003) have shown that fine-grained resuspended sediment acts as seed crystals for the slow precipitation of CaCO 3 and that much of the precipitation actually takes place on dispersed suspended sediment outside of the whitings. They also 88

108 89 observed that precipitation ceases at a supersaturation close to twice that of aragonite. This occurs both in bank waters and in laboratory experiments using resuspended fine grained GBB sediments. Interestingly and perhaps coincidently, this is close to the solubility of the monoclinic form of calcium carbonate, the mineral vaterite (Morse and Mackenzie, 1990). A summary of the maximum precipitation in the GBB throughout the years (with measurements made over several years) is shown in Fig The greatest loss of CaCO 3 (s) (~600 μmol kg -1 ) occurs in the summer months. This is when the waters have the highest salinity (see Fig. 5.2). The precipitated CaCO 3 (s) levels off at high salinity since these older waters are in near equilibrium with aragonite. 700 Bahama Banks CaCO 3 Precipitated (µmol kg -1 ) Month Figure 5.1. The maximum CaCO 3 (s) precipitated on the Grand Bahama Bank over the time of the year. The values are based on the changes in the total alkalinity corrected for salinity.

109 90 Bahama Banks CaCO 3 Loss (µmol kg -1 ) Salinity Figure 5.2. The maximum CaCO 3 (s) precipitated from waters on the Grand Bahama Bank as a function of the salinity. The changes in ph and total carbon dioxide (TCO 2 ), and 49 μm, respectively, of the adjacent waters over the last six decades are in reasonable agreement with the expected values due to the change in the partial pressure of carbon dioxide (pco 2 ) of the atmosphere (80 μatm). To the best of our knowledge, similar studies have not been conducted on the nearby northern Little Bahama Bank (LBB) where space shuttle photographs (Fig. 5.3) indicate the presence of extensive whitings north of the Grand Bahama Island. Two research cruises were made to the LBB, one in July 2003 and the other in May On the first cruise 29 whitings were observed and 6 were sampled, and on the second cruise 43 whitings were sighted and 14 were sampled. Locations of sampled whitings and

110 91 stations are shown in Figs. 5.4 to 5.7. Measurements were made of the CO 2 parameters (ph, TA, TCO 2, pco 2 ), and suspended calcium carbonate was collected on these cruises. Our results, reported in Bustos-Serrano et al., (2009), demonstrate for the first time active precipitation of CaCO 3 (s) in water on the Little Bahama Bank (LBB). Figure 5.3. A photograph (Roll 719; Frame 29) of whitings on the Little Bahama Bank from the NASA Space Shuttle. Photograph courtesy of the Image Science & Analysis Laboratory, NASA Johnson Space Center (

111 Experimental methods General information. Two cruises were made in the Little Bahamas Bank in 2003 and The hydrographic and carbonate measurements made on the cruises are given in Tables 5.1 and 5.2 along with station locations. Seawater samples were collected from the flowing seawater line on the ship. Salinity and temperature of the water were recorded continuously on the cruise. The salinity was also determined on the collected sample using a Guildline Autosal 8410A salinometer. All samples collected in the whitings were filtered through a 0.45 μm filtering system to remove any solid CaCO 3 (s) in the sample. The filtered samples were sealed in plastic bags for analysis in the laboratory. The sediments were collected by hand and were stored in plastic bottles. The 14 C measurements on the whiting material and sediments were measured at Beta Analytical Inc. The Mg and Sr measurements and stable isotope measurements of oxygen and carbon were sent to a specialized laboratory at Texas A&M University managed by Dr. Morse. The carbonate parameters measured on the ship include ph, pco 2, TA and TCO 2. The methods used are briefly discussed below Total alkalinity. The TA was determined using a titration system described in Millero et al., (1993a). The titration system is automatically controlled with a personal computer and the temperature of the measurements (25 ± 0.05 C) was controlled with a constant temperature bath (Neslab, model RTE 221). The TA titrations were analyzed using the

112 93 TFM software which collects the titration data and calculates the ph, E* the standard potential, TA and TCO 2 of the sample (see Chapter 2, sections and 2.1.3). The accuracy of the TA and TCO 2 measurements were checked using Gulf Stream surface seawater and Certified Reference Material (CRM, Batches 57 and 66) provided by Dr. Andrew Dickson, Marine Physical Laboratory, La Jolla, California, USA. The reproducibility of the measurements on the standards was within ± 2 µmol kg -1 for TA and ± 4 µmol kg -1 for TCO ph. The ph was determined by potentiometry as described in Chapter 2, section and using the spectrophotometric procedure of Clayton and Byrne (1993). The measurements were all converted (Lee and Millero, 1995) to the seawater ph SWS scale ([H + ] SWS = [H + ] + [HSO - 4 ] + [HF]). The ph determined by the two methods agrees to within Since the spectroscopic method is more precise than the electrode method, we have used the spectroscopic value of ph for all of our calculations. The m-cresol purple (mcp) indicator was used. The system is fully automated and makes measurements every twelve minutes on a sample of seawater from the flowing seawater line. A microprocessor controlled a KLOEHN syringe and a sampling valve aspirates and injects the seawater sample into the 10 cm optical cell. The software permits five minutes for temperature stabilization. A circulating temperature bath (Neslab RTE-221) regulates the temperature of the sample at 25 ± 0.01 ºC. An Agilent UV/VIS spectrophotometer measures background absorbance of the sample. A Guildline 9540 digital platinum resistance thermometer measured the temperature of the sample.

113 Total inorganic carbon. TCO 2 was determined by coulometry. The measurements were made on the CO 2 stripped from the solution after the addition of H 3 PO 4. The SOMMA (Single-Operator Multiparameter Metabolic Analyzer) system (Johnson et al., 1993) was used in this study on discrete samples. The system was calibrated using CRM throughout the study. The measurement on CRM agreed to ± 2 µmol kg -1 with the assigned value Partial pressure of CO 2 The pco 2 was determined in seawater and the atmosphere using a flowing system similar to the one designed by Wanninkhof and Thoning (1993). The equilibrator used during the cruises was based on the design by Weiss (1981) and modified by Goyet and Peltzer (1993). The fraction of CO 2 in the equilibrated air sample was measured using differential, non-dispersive LI-COR infrared CO 2 /H 2 O analyzer. Samples were measured wet and the signal was corrected for dry air at 1 atm pressure using the measurements from the LI-COR water channel and pressure gauge. The water channel was calibrated using the LI-COR Dew Point Generator, model LI-610, as described in the LI-6262 instrument operation manual. The reference chamber of the IR analyzer was continuously purged with dry, CO 2 free air. This allows the operation of the instrument in the absolute mode. The system was calibrated using three standard gases 294, 324 and 549 ppm in 2003 and 291, 370 and 469 ppm in 2005 (CO 2 in air). The results are reliable to ± 3 µatm. All the calculations of carbonate parameters were made using the CO2sys software (Lewis and Wallace, 1998) and the carbonic acid constants of Millero et al.,

114 95 (2006). Internal consistency studies with input of TA-TCO 2 yielded average values of ph ± and pco 2 of ± µatm compared to the measured values Calcium carbonate and water mass age. Surface waters were collected and filtered on 0.45 μm Nucleopore filters to obtain suspended calcium carbonate solid samples. They were given a quick rinse with calcite-equilibrated pure water to remove adhering seawater and dried. Samples for the determination of suspended calcium carbonate concentrations were dissolved in HCl and the resulting calcium concentration was measured by Inductively Coupled Plasma (ICP) analysis at Texas A&M University by Dr. Morse. Sr/Ca and Mg/Ca ratios on suspended carbonates and bottom sediments were also determined by ICP analyses. Stable isotope ratios were measured by mass spectrometry relative to the PDB standard. The determinations of 14 C for the samples were made by Beta Analytical Inc. The water mass age was estimated from salinity using the 14 C technique-based relationship between salinity and water mass age found on the GBB by Morse et al., (1984). This estimation is based on the assumption that net evaporation rates are similar on the two banks given their close geographic proximity. 5.3 Results. Measurements made on the two cruises have been summarized in a number of contour diagrams of the measured properties (Fig. 5.4 to Fig. 5.7). The contour diagrams of the salinity of the waters on the two cruises are shown in Fig The highest salinities are at the interior due to evaporation over the year and are the oldest waters in the LBB with an estimated residence time of 144 days.

115 Figure 5.4. Contours of the surface salinity of waters in the Little Bahama Bank from cruise in July 2003 and May Blue circles are sampling stations. Interpolation using DIVA from Ocean Data View software. 96

116 Figure 5.5. Contours of the normalized total alkalinity of waters in the Little Bahama Bank from cruises in July 2003 and May

117 Figure 5.6. Contours of the normalized total carbon dioxide of waters in the Little Bahama Bank from cruises in July 2003 and May

118 99 The major changes in the carbonate parameters occur in waters of high salinity as shown in contour plots of the salinity normalized TA (NTA = TA*35/S) (Fig. 5.5) and NTCO 2 (NTCO 2 =TCO 2 *35/S ) (Fig. 5.6), and ph (Fig. 5.7). The values of NTA, for example, are a nearly linear function of salinity as shown in Fig This is due to the increase in the precipitation of CaCO 3 (s) with time. It is also reflected in a decreasing calculated supersaturation state (Ω Aragonite ) of the waters with respect to aragonite (Fig. 5.9). The saturation state of aragonite is defined by Ω Aragonite = [Ca 2+ ][CO 2-3 ]/Ksp * (5.1) where the brackets refer to the concentration (mol kg -1 ) of Ca 2+ and CO 2-3 in solution and K*sp is the solubility of aragonite at the temperature, salinity and pressure of the sample. The variations in NTA across the whitings are clearly evident in the sampling transect plot shown in Fig The dotted line denotes the NTA values for Gulf Stream waters (~2290 µmol kg -1 ). The maximum decrease of NTA is ~470 µmol kg -1 and ~355 µmol kg -1, respectively in 2003 and The large decreases in NTA are in regions where whitings were detected and indicate that active precipitation was occurring. The values of NTCO 2 (Fig. 5.11) also show much lower values than the values for the Gulf Stream (~1960 µmol kg -1 ). The maximum decreases in NTCO 2 are ~490 µmol kg -1 and ~360 µmol kg -1, respectively in 2003 and The changes in the pco 2 (μatm) shown in Fig are quite large when active precipitation occurs (see Fig. 5.13). Due to delays in the measurements of pco 2 on the ship, it was not possible to match up the change with those determined for TA and TCO 2. The maximum increases in pco 2 are ~ 135 μatm for both years. It should be pointed out that the calculated values of pco 2 in 2003 are higher by ~80 μatm than the 2005 results due to changes in temperature and

119 100 differences in TCO 2. These differences agree with the calculated value using TA and TCO 2. Figure 5.7. ph contours in the Little Bahama Bank from cruises in 2003 and 2005.

120 101 Little Bahama Bank NTA (µmol kg -1 ) Salinity Figure 5.8. Values of NTA plotted against salinity which is a proxy for water residence time with the highest salinity occurring at about 144 d Figure 5.9. Contours of the saturation state of aragonite of waters in the Little Bahama Bank from cruises in 2003 and 2005.

121 102 Results from the two cruises to the LBB are given in Table 5.1 for July 2003 and in Table 5.2 for May The suspended solid CaCO 3 (s) in the waters with and without whitings have also been measured (Fig. 5.13). The whitings contain ~10 times more CaCO 3 (s) than the non-whiting waters. The heaviest whitings could be produced entirely by local precipitation. 14 C analysis of calcium carbonate filtered from the whitings gave an average age of about 700 years and the fine fraction of the sediment had a similar average age of about 1000 years. This indicates that a dominant fraction of the calcium carbonate in the whiting is not freshly precipitated, but rather resuspended sediment. As shown in Fig. 5.14, the Mg and Sr content of the whiting material and the sediments in LBB are similar as those found in the GBB (Morse et al., 2003). The stable isotopes ratios of oxygen (δ 18 O) and carbon (δ 13 C) of the sediments in the LBB and GBB shown in Fig are also similar.

122 103 Table 5.1. Carbonate measurements made on July 2003 Bahamas Cruise. Bahamas Cruise Station Latitude Longitude Salinity Temp ph TA TCO 2 Deg Deg

123 104 Station Latitude Longitude Salinity Temp ph TA TCO 2 Deg Deg

124 105 Table 5.2. Carbonate measurements made on May 2005 Bahamas Cruise. Bahamas Cruise Station Latitude Longitude Salinity Temp ph TA TCO 2 Deg Deg

125 106 Station Latitude Longitude Salinity Temp ph TA TCO

126 Seawater 2200 NTA µmo kg Whitings 2003 Cruise Seawater NTA µmo kg Whitings 2005 Cruise Station Figure The values of NTA plotted for the stations occupied on the cruises.

127 Seawater NTCO 2 µmo kg Whitings 2003 Cruise Seawater 1900 NTCO 2 µmo kg Whitings 2005 Cruise Station Figure The values of NTCO 2 plotted for the stations occupied on the cruises.

128 Figure The values of pco 2 plotted for the stations occupied on the cruises. Reference line at 350 µatm is the atmospheric value. 109

129 110 8 Number of Observations Open Water (mg CaCO 3 L -1 ) Number of Observations (1.5 mg) Whiting Water (mg CaCO 3 L -1 ) Figure The amount of suspended CaCO 3 (s) in whiting and non whiting waters. Note the difference in the scales between open ocean and whitings waters.

130 Sr/Ca Mg/Ca Figure The Sr/Ca versus the Mg/Ca ratios of whiting and fine-grained bottom LBB sediment. Whiting values are solid circles and bottom sediments are open circles GBB Whiting Sediment GBB Bottom Sediment δ 18 O %o Figure δ 13 C versus δ 18 O, relative to the PDB standard, of whiting and bottom sediments. Whiting values are solid circles and bottom sediments are open circles.

131 Discussion. There are many major similarities between the waters on the LBB and GBB associated with whitings. In both cases, the oldest most saline waters are trapped against the major islands, northern Grand Bahama Island on the LBB and western Andros Island on the GBB. Major losses of NTA occur across both banks indicating calcium carbonate precipitation over broad areas resulting in a major decrease in the saturation state of the waters with respect to aragonite. The residence time of the water on the LBB is about half that on the GBB, and the relative changes in water chemistry are roughly proportional to this difference in residence time. Both, Broecker and Takahashi (1966) and Morse et al., (1984) used changes in Δ 14 C in the waters, which were closely correlated with salinity, to obtain residence time distributions of Bank waters. From these sets of observations, it is possible to make estimates of precipitation rates as a function of supersaturation. Certainly one of the most important observations is that the calcium carbonate in the LBB is similar in age to the underlying sediments as is the case for GBB whiting carbonates. This indicates that on both banks, the formation of whitings is primarily due to the resuspension of fine-grained sedimentary carbonates. Thus, whitings are not brought about by some biologically induced pseudo-homogenous nucleation process as has often been argued for the GBB with little-to-no supporting evidence (Robbins and Blackwelder 1992). It is likely that the primary precipitation process for calcium carbonate on the LBB is by epitaxial growth on the resuspended fine-grained sediment which is predominantly small aragonite needles as is the case for the GBB (Morse et al., 2003).

132 113 Although there was considerable variability in the data and it may not give statistically strong average values, suspended calcium carbonate average concentrations in and out of whitings were about the same on the LBB and the northern GBB west of Andros Island in the region where whitings were observed (~5 mg L -1 inside and ~0.5 mg L -1 outside). It should be noted that Shinn et al., (1989) found average values about twice as high in their study of the GBB whitings (~10 mg L -1 inside and ~1.3 mg L -1 outside) perhaps reflecting differing conditions at that time. Although there were several major similarities, there were also some important differences observed with the whitings on the LBB compared with those on the GBB. Visually whitings on LBB often appeared much denser and had a roiling appearance that wasn t observed on the GBB, even though the suspended calcium carbonate concentrations were similar as noted in the previous paragraph. This was also reflected in the major differences in the NTA and pco 2 values of the waters inside and outside of the whitings. These changes in whiting waters were roughly enough to produce the observed concentration changes in suspended calcium carbonate. However, we do not think this is a true quantitative relationship due to the relatively rapid resuspension and settling of particles that takes place in whitings (see Shinn et al., 1989). A possible explanation for the fact that we did not observe such changes in water chemistry in GBB whitings is that either the duration of a whiting event is longer on the LBB, or that the whiting and associated water packet are more tightly coupled on the LBB.

133 Synopsis. This first study of the LBB shown in satellite photographs (Fig. 5.16) and the site seen from the ship (Fig and Fig. 5.18) show the presence of whitings. The site reveals that the general hydrography, variations in carbonate-associated water chemistry, and characteristics of the suspended calcium carbonate were remarkably similar to those on the well studied GBB. For both banks, the observations support the idea that the primary process responsible for the observed change in water chemistry is the precipitation of calcium carbonate on resuspended carbonate sediments. On the LBB major changes in water chemistry inside the whitings versus the surrounding water, indicate calcium carbonate precipitation in the whitings, was clearly observed for the first time. Figure Little Bahama Bank (pale blue) just north of the Great Bahama Island show some whitings using satellite imagery on Google Earth (Oct. 25, 2008).

134 115 Figure Cruise Track at LBB cruise in May 2005 (blue line) some whitings were observed from the R/V Walton Smith bridge. Blue dots reflect the size of whitings, most of them with irregular shapes.

135 Figure Whiting observed at the LBB cruise in July 2003 on the bridge of the R/V Walton Smith, University of Miami. 116