1. The Seawater CO 2 -Carbonate System Close-up of Oolitic limestone http://en.wikipedia.org/wiki/File:CarmelOoids.jpg Carmel formation, Utah, Jurassic

2. Some of the major geochemical roles of the CO 2 system in seawater include: seawater pH control and buffering source of carbon for photosynthesis long-term sink for carbon via carbonate precipitation and subsequent burial and preservation of limestone and dolomite formation of carbonate reefs exchange of CO 2 with the atmosphere: CO 2 is a major greenhouse gas. Oceans are both a source and sink for atmospheric CO 2 depending on location; a net sink overall. source of biogenic carbonates that are important paleoindicators for a variety of parameters

3. Use of sign inspired by talk given by Andrew Dickson, a noted CO2-system chemist from Scripps. The Seawater CO 2 -Carbonate System

4. The carbonate system is one of the most important chemical and biogeochemical systems on earth. CO 2 (g)  CO 2 (aq) + H 2 O H 2 CO 3 H + + HCO 3 - H + + CO 3 2- K o K 1 K 2 For seawater with a salinity of 35 and a temperature of 25 o C: pK o = 1.547 pK 1 = 5.847 pK 2 = 8.915 Air Sea Proton #1Proton #2 (From Millero, Table 7.4) O=C=O CO 2 O=C-OH O - HCO 3 - Structures of CO 2 gas and bicarbonate ion

5. From, Millero, Chemical Oceanography, 1996. p 246 H 2 CO 3 H 2 O + CO 2 The total CO 2 stays constant but the speciation depends on pH Equivalence points pH = pK 1 pH = pK 2

6. Terminology related to the CO 2 system in seawater DIC - Dissolved inorganic carbon (CO 2(g) +H 2 CO 3 +HCO 3 - +CO 3 2- ) CO 2 (g) : gaseous CO 2 CO 2 (aq) : gaseous CO 2 that is dissolved in water Carbonate ion: CO 3 2- Bicarbonate ion: HCO 3 - Carbonic acid: H 2 CO 3 (non-charged,neutral species) Total CO 2 (  CO 2 ): Sum of all dissolved components of inorganic carbon, including CO 2 (aq), H 2 CO 3, HCO 3 -, and CO 3 2-. Total CO 2 = DIC PIC – Particulate inorganic carbon (calcite & aragonite minerals)

7. Distribution of DIC in the ocean- Vertical and horizontal distributions Total DIC (=Total CO 2 ) is lowest (but not zero!) in surface waters and is enriched in deeper water - the enrichment is greatest in the deep Pacific due to water mass age. Why is this pattern observed?

8. Dissolution of CO 2 in water results in formation of carbonic acid, which dissociates to yield bicarbonate and carbonate plus protons. Thus, the addition of CO 2 to water increases the {H + } and therefore lowers the pH of the solution. Conversely, removal of CO 2 from solution removes {H + } and increases the pH CO 2 (aq) + H 2 O H 2 CO 3 H + + HCO 3 - H + + CO 3 2- pH in Seawater – Complex control by CO 2 system and Alkalinity

9. Biological uptake of carbon by marine plants is mainly as CO 2(g) or H 2 CO 3 i.e. neutral species. (some phytoplankton can take up HCO 3 - and convert it to CO2 via carbonic anhydrase) The uptake of CO 2 (g) by photosynthetic organisms (or chemosynthetic organisms) will raise the pH of the system due to shift in the equilibria to the left (in the direction consuming H +. Conversely, respiration of organic matter reverses the cycle and liberates CO 2 which will dissociate and lower the pH (increase {H + }). CO 2 (aq) + H 2 O H 2 CO 3 H + + HCO 3 - H + + CO 3 2- Remove CO 2, consume H + and raise pH Add CO 2, add H + and lower pH

10. Vertical distribution of pH in the ocean 2 4 6 0 Depth (km) 7.57.77.98.1 pH Atlantic Indian Pacific There is an ocean- wide pH minimum starting just below the euphotic zone and extending to 500-1000 m. pH is lower in deep Pacific than in Atlantic - due to water mass age and accumulation of respired CO 2 !

11. Seawater alkalinity - a measure of the buffering capacity Simply put: “The amount of negative charge in seawater that is able to accept a proton (hydrogen ions) during the titration of seawater with strong acid to the point where essentially all the carbonate species are protonated” (paraphrased from Pilson, p 114). Alkalinity is not the total negative charge in solution, but rather just the concentration of negatively charged species that will accept H + above certain pH end-point - defined by the method of titration - usually around pH 3.5 - 4.5 Chloride is a negatively charged species, but Cl - will not accept a proton in aqueous solution, even at pH 0!

12. For most natural waters, total alkalinity (TA) can be simplified to: (TA) = [HCO 3 - ] + 2 [CO 3 2- ] + [B(OH) 4 - ] + [OH - ] - [H + ] ~95% of the alkalinity in seawater is comprised of the carbonate alkalinity; Carbonate Alkalinity (CA) = [HCO 3 - ] + 2 [CO 3 2- ] Carbonate Alkalinity (in molal units) is always greater than Total CO 2 (in molal units) because each unit of CO 3 2- contributes 2 units of alkalinity (can accept 2 protons) Borate contributes about 5% to the alkalinity and needs to be taken into account.

13. Alkalinity according to Dickson (1992) (a detailed view)  Alkalinity = [HCO 3 - ] + 2[CO 3 2- ] + [B(OH) 4 - ] + [OH - ] + [HPO 4 2- ] + 2[PO 4 3- ] + [SiO(OH) 3 - ] + [HS - ] + [NH 3 ] + [all other unidentified weak bases] - [H + ] - [HSO 4 - ] - [HF] - [H 3 PO 4 ] - [all other unidentified acids] Alkalinity is not strictly related to pH. For example: Deep ocean water has higher alkalinity than the surface - but a lower pH (higher acidity).

14. Factors affecting Alkalinity and Total CO 2 in seawater  Alkalinity is not affected by T and P (because it is a charge balance).  Alkalinity increases with dissolution of carbonate minerals (which release HCO 3 - or CO 3 2- ). Dissolution of CaCO 3 releases CO 3 2-, thereby increasing alkalinity. Likewise, precipitation of carbonate minerals consumes (decreases) alkalinity. Carbonate precipitation also affects  CO 2.  Photosynthesis and respiration consume and add CO 2 respectively, but do not affect alkalinity. This is because release of CO 2 and subsequent hydration and dissociation yields HCO 3 - + H + (one unit of H + for every unit of negative charge alkalinity). Remineralization does however, increase  CO 2.  The exception to this rule is respiration with sulfate as the electron acceptor. Sulfate reduction generates HS - which increases alkalinity. The CO 2 generated in sulfate reduction only increases  CO 2 and not alkalinity.

15. Vertical and horizontal distribution of Alkalinity in the ocean Similar to that of total DIC - low in surface waters, increasing with depth in thermocline. Higher in deep Pacific than in Atlantic - due to water mass age and inputs of CO 3 2- from CaCO 3 dissolution Fig. 15.10 in Libes DIC Alkalinity

16. Precipitation and dissolution of carbonate minerals in the ocean All seawater contains the ions Ca 2+, CO 3 2- and HCO 3 -. The effective concentrations (i.e. activities) of these species, together with the pH, temperature, pressure and ionic strength determine whether the solution is saturated or undersaturated with respect to CaCO 3 minerals. Cocolithophore Emiliania huxleyi a haptophyte phytoplankter secretes plates (liths) of calcite (CaCO 3 )

17. Carbonate Minerals Calcite CaCO 3 Ca 2+ + CO 3 2- K sp = 4.47 x 10 -9 @ 25 o C and Ionic strength of 0 (a std condition) Aragonite CaCO 3 Ca 2+ + CO 3 2- K sp = 6.02 x10 -9 @ 25 o C and Ionic strength of 0 (a std condition) Aragonite has the larger K sp, therefore it is more soluble. Aragonite is more amorphous (less ordered crystal) and is more soluble than calcite. These two compounds differ only in their crystalline structure not their chemical formula which is CaCO 3 in both cases.

18. Calcite is the predominant form of CaCO 3 in the ocean and it is more stable than aragonite (amorphous CaCO 3 ). The organisms that precipitate calcite include (Cocolithophores and foraminifera). Organisms that precipitate aragonite include Corals & Pteropods. Coccolithophore Foraminiferan Pteropod Dinoflagellate cyst

19. Precipitation of biogenic carbonates and their subsequent burial in marine sediments represents the single largest export of carbon from the biosphere. Biogenic carbonates in sedimentary rocks (e.g. limestones and dolomites) are the single largest reservoir of carbon on Earth. Most of this carbonate is derived from planktonic microorganisms Not all biogenic carbonate is preserved in sediments - much dissolves in the deep sea. The precipitation and dissolution of CaCO 3 depends on the physicochemical conditions in seawater There is 1400 times more Carbon tied up in carbonate rocks than there is in DIC in the ocean!

20. Nearly all surface ocean waters are supersaturated with respect to calcite and aragonite; deep waters are undersaturated.

21. Despite surface supersaturation, spontaneous precipitation of calcite or aragonite in surface waters does not occur (except at very high pH's) due partly to interaction of Mg 2+ with CaCO 3 crystal surfaces. Only in very warm, saline waters where CO 2 solubility is low (hence  CO 2 is low) will CaCO 3 ppt out as aragonite without biocatalysis. Carbonate Ooids are examples of spontaneously precipitated carbonates – currently found on Bahamas platform – but extensive geological deposits exist Calcifying organisms overcome the Mg 2+ problem with enzymes and intracellular compartmentalization of pH etc. http://www.iun.edu/~geos/

22. Calcification: Calcium carbonate precipitation can be written simply as: Ca 2+ + CO 3 2-   CaCO 3 (s) But biogenic CaCO 3 precipitation appears to occur primarily by the following reaction mechanism: Ca 2+ + 2HCO 3 -   CaCO 3(s) + CO 2 + H 2 O Thus, per mole of CaCO 3 formed, calcification i) consumes 2 mole of alkalinity, ii) consumes 1 mole of DIC and iii) produces 1 mole of CO 2 (i.e. increases pCO 2 )

23. The K eq values for the CO 2 system reactions are a function of temperature & pressure therefore so is CaCO 3 solubility As Temp goes down, pH goes down; K sp of CaCO 3 goes up (more soluble) (retrograde solubility) As Pressure goes up, pH goes down; K sp of CaCO 3 goes up (more soluble) These effects are due the fact that CO 2 gas and charged vs. neutral species are involved in the equilibrium. Gases are more soluble at higher pressures and lower temperatures, favoring CO 2 (g) dissolution, hence more carbonic acid forms. Also, as pressure increases, formation of charged species is favored because the ions have a lower partial molal volume than the solid (or neutral species) due to electrostriction.

24. The degree of saturation (Omega) can be expressed as: Omega is given as an output in the CO2SYS program. If IP > K sp *, then solution is supersaturated (  > 1). If IP < K sp * then solution is undersaturated (  < 1). Solubility of CaCO 3 depends mostly on variations in CO 3 2- rather than Ca 2+ because Ca 2+ is nearly constant in the ocean. The rate of dissolution of CaCO 3 is an exponential function of the degree of undersaturation.

25. CaCO 3 is not found in the surficial sediments in the deepest parts of the sea (> ~5000 m) to any great extent for at least two reasons. 1) The solubility of CaCO 3 increases as Pressure  and as Temp.  2) pH decreases with depth and more CaCO 3 will dissolve. The lysocline is the depth at which significant dissolution of calcite begins. This depth is different for different ocean water masses. The CCD (Calcite compensation depth) is the depth at which the dissolution of CaCO 3 minerals equals the supply rate (rain rate). No significant accumulation of CaCO 3 occurs below this depth.

26. Emerson and Hedges Fig 12.12 Δ

27. Source: Open University: Ocean chemistry and deep sea sediments Places where CaCO 3 dominates the sediments are relatively shallow (< 5000 m) CaCO 3

28. The CCD for aragonite is much shallower than for calcite because aragonite is more soluble (larger K sp ). The CCD for calcite is shallower in the Pacific (3.5 km) than in the Atlantic (5 km) due to the lower pH of the Pacific deep waters (caused by age and CO 2 production from respiration). Many factors govern the CCD including the rate of supply, chemical composition, minerology, size and shape, rate of bioturbation. Larger particles may not dissolve quickly. The distribution of CaCO 3 oozes as they are called (sediments with > 75% CaCO 3 ) is largely restricted to shallower parts of the oceans (see next slide figure from Open University text).

30. Calculation of all the parameters of the CO 2 system in seawater using the CO2SYS program (available for free download at: http://cdiac.esd.ornl.gov/oceans/co2rprt.html). CO2SYS will do all the calculations for you provided you have input data for two of the four main parameters of the CO 2 system: Total AlkalinitypCO 2 pHTotal CO 2 With input of two parameters (plus temperature, salinity, pressure, silicate and phosphate data) the other two parameters of the CO 2 system will be predicted as well as the concentration of various species, the degree of calcite or aragonite saturation, and more. The program is very easy to use.

31. pH values on the total scale (pH tot ) are about:.09 units lower than those on the free scale,.01 units higher than those on the seawater scale, and.13 units lower than those on the NBS scale. pH Scales (defined in CO2SYS) pH NBS ( National Bureau of Standards ; standard lab pH buffers are NBS, but they are low ionic strength and not great for SW. pH seawater pH total pH free Dickson recommends these Differences in these scales have to do with how they consider the sulfuric acid and hydrofluoric acid components of seawater

32. Exchange of CO 2 (g) between the atmosphere and the ocean Portions of the ocean surface are super saturated with CO 2(g) while other portions are undersaturated. Only the CO 2 (g) part of the total CO 2 system can exchange with the atmosphere. Thus, knowledge of the partial pressure of carbon dioxide (pCO 2 ) is critical for understanding exchanges of carbon between the atmosphere and oceans. There is about 50 times more Total CO 2 dissolved in the oceans than there is CO 2 in the atmosphere.

33. CO 2 (aq) is in equilibrium with the atmosphere such that: [CO 2 aq ] = H CO2 * pCO 2 Where H CO2 is the Henry’s Law constant for CO 2 and pCO 2 is the partial pressure of CO 2. The Henry’s Law constant is essentially the equilibrium constant for the dissolution of the gas: CO 2 (g) CO 2 (aq) K eq = {CO 2 (aq )}/ {CO 2 (g) } The activity of CO 2 in the gas is essentially its partial pressure

34. The current concentration of CO 2 in the atmosphere (in 2012) is about 392 ppm, or 0.0392%. This concentration has already increased 40% from pre-industrial values and is expected to nearly double in the next century. Source: C. D. Keeling http://cdiac.esd.ornl.gov/trends/co2/sio-mlo.htm Atmospheric CO 2 is increasing dramatically The implications for marine systems are huge! Ron enters graduate school @ 340 ppm Ron born at 318 ppm 392 in Sep 2012 The increase is 78 ppm in 54 years – a 25% increase Get the latest CO 2 concentration at http://co2now.org/

35. Source: Buddemeir et al. Pew Report on Coral Reefs and Global Climate Change

36. Consequences of global increase of CO 2 in atmosphere t Greenhouse warming t Sea level rise - polar ice decline t Enhanced terrestrial primary productivity t Decreased seawater pH (more carbonic acid)  changes in phytoplankton physiology/ecology  decreased calcification by corals and other marine organisms t decrease of pH in rain/snow - greater terrestrial weathering t Indirect effects - many

37. 26% increase for pH 8.2  8.1 58% increase for pH 8.2  8.0

38. Source: Buddemeir et al. Pew Report on Coral Reefs and Global Climate Change 2x CO 2

39. The oceans do not always achieve equilibrium with respect to atmospheric CO 2 Only about half of the CO 2 input to the atmosphere by Man’s activities since the dawn of the industrial age has accumulated in the atmosphere. The other half has been absorbed by either the oceans or the terrestrial biota. The ocean’s response takes time. The marine biota add or remove CO 2 in surface waters on short time scales, thereby affecting direction of the CO 2 flux. This is due to sluggish kinetics of the equilibria of gas exchange and the fact that the ocean is a layered system with a relatively long residence- and mixing time.

40. Out of the ocean Into the ocean

41. Fish otoliths (ear stones) are made of aragonite/protein layers Checkley et al. found that growing larval White Sea Bass at elevated pCO 2 caused otoliths to be 8% and 16% larger in the 1000 and 2500 µatm treatments compared to the 430 µatm controls Control

42. Stop !

44. IPCC-FAR

45. In calculating the K sp the activity of the ionic species should be used. In practice, marine chemists would measure the concentration of Ca 2+ and CO 3 2- at which precipitation occurs. The resulting solubility product would be the apparent K sp ’ or stoichiometric constant. It’s value would depend on the conditions such as temp, pressure, and Ionic strength.

46. In practice alkalinity is measured by titration. The amount of H + in equivalents per kg needed to titrate 1 kg of seawater to the bicarbonate/H 2 CO 3 equivalence point. Modern methods involve coulombic titrations to determine end point.

47. Depth profiles of carbonate mineral saturation state in the Atlantic and Pacific Oceans. An Omega value of 1 indicates saturation; above 1 is supersaturated; below, undersaturated. (From Millero, Chemical Oceanography, 1996. pp 274 & 275.

49. HBa  H + + Ba - H 2 O  H + + OH - Equilibria Since there are no other cations to balance the [Ba - ], H + will adjust to equal the concentration of Ba - to satisfy both equilibria. This sets pH at 4 in this case

51. Equivalence points – where pH = pK Consider dissociation of the weak acid: H 2 CO 3  H + + HCO 3 - K a = {H + } {HCO 3 - }/{H 2 CO 3 } K a {H 2 CO 3 } / {HCO 3 - } = {H + } Rearrange to isolate {H + } pK a {H 2 CO 3 } / {HCO 3 - } = pH Take “p” or negative Log of both sides {H 2 CO 3 } / {HCO 3 - } = pH/ pK a At pH = pK a, then {H 2 CO 3 } / {HCO 3 - } = 1 Thus, at pH =pK a, the {H 2 CO 3 } and {HCO 3 - } must be equal, hence the equivalence point or

53. pH control in sea water - seawater pH varies from 7.9-8.4 with an average between 8.1 and 8.2. Seawater pH is controlled largely by the reaction: HCO 3 - H + + CO 3 2- The equilibrium expression for this reaction is: where K a is the dissociation constant of the bicarbonate. (~10 -8.9 ) i.e. pK a ~8.9 Looking at this equilibrium expression another way, the pH will depend on the ratio of bicarbonate to carbonate and vice versa. Measurement of seawater pH is not straightforward due to the complexity of the solution and its ionic strength (I=0.7). Several different pH scales are in use (depending on buffers used for standards). The NBS pH scale is the most common.

54. Ken Caldeira Climate System Modeling Group Lawrence Livermore National Laboratory 7000 East Avenue, L-103 Livermore, CA 94550, U.S.A. E-mail: kenc@LLNL.gov Robert Berner Department of Geology and Geophysics Yale University New Haven, CT 06520-8109, U.S.A. E-mail: berner@hess.geology.yale.edukenc@LLNL.govberner@hess.geology.yale.edu http://www.sciencemag.org/cgi/con tent/full/286/5447/2043a