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The Seawater CO 2 -Carbonate System

The Seawater CO 2 -Carbonate System. Close-up of Oolitic limestone. http://en.wikipedia.org/wiki/File:CarmelOoids.jpg Carmel formation, Utah, Jurassic. Some of the major geochemical roles of the CO 2 system in seawater include: seawater pH control and buffering

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The Seawater CO 2 -Carbonate System

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  1. The Seawater CO2-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 CO2 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 CO2 with the atmosphere: CO2 is a major greenhouse gas. Oceans are both a source and sink for atmospheric CO2 depending on location; a net sink overall. • source of biogenic carbonates that are important paleoindicators for a variety of parameters

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

  4. Proton #1 Proton #2 O=C=O CO2 • The Seawater CO2-Carbonate System • The carbonate system is one of the most important chemical and biogeochemical systems on earth. • CO2 (g) •  • CO2 (aq) + H2O <=> H2CO3<=> H+ + HCO3- <=> H+ + CO32- • Ko K1 K2 • For seawater with a salinity of 35 and a temperature of 25 oC: • pKo = 1.547 • pK1 = 5.847 • pK2 = 8.915 Air Sea (From Millero, Table 7.4) O=C-OH O- Structures of CO2 gas and bicarbonate ion HCO3-

  5. H2CO3 H2O + CO2 Equivalence points pH = pK1 pH = pK2 The total CO2 stays constant but the speciation depends on pH From, Millero, Chemical Oceanography, 1996. p 246

  6. Terminology related to the CO2 system in seawater DIC- Dissolved inorganic carbon (CO2(g)+H2CO3+HCO3-+CO32-) CO2(g) : gaseous CO2 CO2(aq): gaseous CO2 that is dissolved in water Carbonate ion:CO32- Bicarbonate ion:HCO3- Carbonic acid:H2CO3 (non-charged,neutral species) Total CO2 (CO2): Sum of all dissolved components of inorganic carbon, including CO2(aq), H2CO3, HCO3-, and CO32-. Total CO2 = DIC PIC – Particulate inorganic carbon (calcite & aragonite minerals)

  7. Distribution of DIC in the ocean- Vertical and horizontal distributions Total DIC (=Total CO2) 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. pH in Seawater – Complex control by CO2 system and Alkalinity Dissolution of CO2 in water results in formation of carbonic acid, which dissociates to yield bicarbonate and carbonate plus protons. CO2 (aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + CO32- Thus, the addition of CO2 to water increases the {H+} and therefore lowers the pH of the solution. Conversely, removal of CO2 from solution removes {H+} and increases the pH

  9. Biological uptake of carbon by marine plants is mainly as CO2(g) or H2CO3 i.e. neutral species. (some phytoplankton can take up HCO3- and convert it to CO2 via carbonic anhydrase) The uptake of CO2(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 CO2 which will dissociate and lower the pH (increase {H+}). CO2 (aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + CO32- Remove CO2, consume H+ and raise pH Add CO2, add H+ and lower pH

  10. Vertical distribution of pH in the ocean 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 CO2! pH 7.5 7.7 7.9 8.1 0 2 Depth (km) Atlantic 4 Pacific 6 Indian

  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) = [HCO3-] + 2 [CO32-] + [B(OH)4-] + [OH-] - [H+] ~95% of the alkalinity in seawater is comprised of the carbonate alkalinity; Carbonate Alkalinity (CA) = [HCO3-] + 2 [CO32-] Carbonate Alkalinity (in molal units) is always greater than Total CO2 (in molal units) because each unit of CO32- 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 = [HCO3-] + 2[CO32-] + [B(OH)4-] + [OH-] + [HPO42-] + 2[PO43-] + [SiO(OH)3-] + [HS-] + [NH3] + [all other unidentified weak bases] - [H+] - [HSO4-] - [HF] - [H3PO4] - [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 CO2 in seawater • Alkalinity is not affected by T and P (because it is a charge balance). • Alkalinityincreaseswith dissolution of carbonate minerals (which release HCO3- or CO32-). Dissolution of CaCO3 releases CO32-, thereby increasing alkalinity. Likewise, precipitation of carbonate minerals consumes (decreases) alkalinity. Carbonate precipitation also affects CO2. • Photosynthesis and respiration consume and add CO2 respectively, but do not affect alkalinity. This is because release of CO2 and subsequent hydration and dissociation yields HCO3- + H+ (one unit of H+ for every unit of negative charge alkalinity). Remineralization does however, increase CO2. • The exception to this rule is respiration with sulfate as the electron acceptor. Sulfate reduction generates HS- which increases alkalinity. The CO2 generated in sulfate reduction only increases CO2 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. DIC Alkalinity Higher in deep Pacific than in Atlantic - due to water mass age and inputs of CO32- from CaCO3 dissolution Fig. 15.10 in Libes

  16. Precipitation and dissolution of carbonate minerals in the ocean All seawater contains the ions Ca2+, CO32- and HCO3-. 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 CaCO3 minerals. Cocolithophore Emiliania huxleyi a haptophyte phytoplankter secretes plates (liths) of calcite (CaCO3)

  17. Carbonate Minerals Calcite CaCO3 <=> Ca2+ + CO32-Ksp = 4.47 x 10-9 @ 25oC and Ionic strength of 0 (a std condition) Aragonite CaCO3 <=> Ca2+ + CO32-Ksp = 6.02 x10-9 @ 25oC and Ionic strength of 0 (a std condition) Aragonite has the larger Ksp, 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 CaCO3 in both cases.

  18. Calcite is the predominant form of CaCO3 in the ocean and it is more stable than aragonite(amorphous CaCO3). The organisms that precipitate calcite include (Cocolithophores and foraminifera). Organisms that precipitatearagonite include Corals & Pteropods. Coccolithophore Foraminiferan Dinoflagellate cyst Pteropod

  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 There is 1400 times more Carbon tied up in carbonate rocks than there is in DIC in the ocean! Not all biogenic carbonate is preserved in sediments - much dissolves in the deep sea. The precipitation and dissolution of CaCO3 depends on the physicochemical conditions in seawater

  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 Mg2+ with CaCO3 crystal surfaces. • Only in very warm, saline waters where CO2 solubility is low (hence CO2 is low) will CaCO3 ppt out as aragonite without biocatalysis. • Carbonate Ooids are examples of spontaneously precipitated carbonates – currently found on Bahamas platform – but extensive geological deposits exist http://www.iun.edu/~geos/ • Calcifying organisms overcome the Mg2+ problem with enzymes and intracellular compartmentalization of pH etc.

  22. Calcification: Calcium carbonate precipitation can be written simply as: Ca2+ + CO32-  CaCO3 (s) But biogenic CaCO3 precipitation appears to occur primarily by the following reaction mechanism: Ca2+ + 2HCO3-  CaCO3(s) + CO2 + H2O Thus, per mole of CaCO3 formed, calcification i) consumes 2 mole of alkalinity, ii) consumes 1 mole of DIC and iii) produces 1 mole of CO2 (i.e. increases pCO2)

  23. The Keq values for the CO2 system reactions are a function of temperature & pressure therefore so is CaCO3 solubility As Temp goes down, pH goes down; Ksp of CaCO3 goes up (more soluble) (retrograde solubility) As Pressure goes up, pH goes down; Ksp of CaCO3 goes up (more soluble) These effects are due the fact that CO2 gas and charged vs. neutral species are involved in the equilibrium. Gases are more soluble at higher pressures and lower temperatures, favoring CO2 (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 > Ksp*, then solution is supersaturated ( > 1) . If IP < Ksp* then solution is undersaturated ( < 1) . Solubility of CaCO3 depends mostly on variations in CO32- rather than Ca2+ because Ca2+ is nearly constant in the ocean. The rate of dissolution of CaCO3 is an exponential function of the degree of undersaturation.

  25. CaCO3 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 CaCO3 increases as Pressure  and as Temp.  2) pH decreases with depth and more CaCO3 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 CaCO3 minerals equals the supply rate (rain rate). No significant accumulation of CaCO3 occurs below this depth.

  26. Δ Emerson and Hedges Fig 12.12

  27. Places where CaCO3 dominates the sediments are relatively shallow (< 5000 m) CaCO3 CaCO3 CaCO3 CaCO3 CaCO3 CaCO3 CaCO3 CaCO3 Source: Open University: Ocean chemistry and deep sea sediments

  28. The CCD for aragonite is much shallower than for calcite because aragonite is more soluble (larger Ksp). 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 CO2 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 CaCO3 oozes as they are called (sediments with > 75% CaCO3) is largely restricted to shallower parts of the oceans (see next slide figure from Open University text).

  29. Calculation of all the parameters of the CO2 system in seawater using the CO2SYS program (available for free download at: http://cdiac.esd.ornl.gov/oceans/co2rprt.html). CO2SYSwill do all the calculations for you provided you have input data for two of the four main parameters of the CO2 system: Total AlkalinitypCO2pHTotal CO2 With input of two parameters (plus temperature, salinity, pressure, silicate and phosphate data) the other two parameters of the CO2 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.

  30. pH Scales (defined in CO2SYS) pHNBS (National Bureau of Standards; standard lab pH buffers are NBS, but they are low ionic strength and not great for SW. pHseawater pHtotal pHfree Differences in these scales have to do with how they consider the sulfuric acid and hydrofluoric acid components of seawater Dickson recommends these pH values on the total scale (pHtot) 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.

  31. Exchange of CO2 (g) between the atmosphere and the ocean • Portions of the ocean surface are super saturated with CO2(g) while other portions are undersaturated. • Only the CO2(g) part of the total CO2 system can exchange with the atmosphere. Thus, knowledge of the partial pressure of carbon dioxide (pCO2) is critical for understanding exchanges of carbon between the atmosphere and oceans. • There is about 50 times more Total CO2 dissolved in the oceans than there is CO2 in the atmosphere.

  32. CO2(aq) is in equilibrium with the atmosphere such that: [CO2aq] = HCO2 * pCO2 Where HCO2is the Henry’s Law constant for CO2 and pCO2 is the partial pressure of CO2. The Henry’s Law constant is essentially the equilibrium constant for the dissolution of the gas: CO2(g) <=> CO2(aq) Keq = {CO2(aq)}/ {CO2(g)} The activity of CO2 in the gas is essentially its partial pressure

  33. The current concentration of CO2 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. Ron enters graduate school @ 340 ppm Atmospheric CO2 is increasing dramatically 392 in Sep 2012 The implications for marine systems are huge! The increase is 78 ppm in 54 years – a 25% increase Ron born at 318 ppm Get the latest CO2 concentration at http://co2now.org/ Source: C. D. Keeling http://cdiac.esd.ornl.gov/trends/co2/sio-mlo.htm

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

  35. Consequences of global increase of CO2 in atmosphere • Greenhouse warming • Sea level rise - polar ice decline • Enhanced terrestrial primary productivity • Decreased seawater pH (more carbonic acid) • changes in phytoplankton physiology/ecology • decreased calcification by corals and other marine organisms • decrease of pH in rain/snow - greater terrestrial weathering • Indirect effects - many

  36. 58% increase for pH 8.2  8.0 26% increase for pH 8.2  8.1

  37. 2x CO2 Source: Buddemeir et al. Pew Report on Coral Reefs and Global Climate Change

  38. The oceans do not always achieve equilibrium with respect to atmospheric CO2 • 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. • The marine biota add or remove CO2 in surface waters on short time scales, thereby affecting direction of the CO2 flux. • Only about half of the CO2 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.

  39. Into the ocean Out of the ocean

  40. Fish otoliths (ear stones) are made of aragonite/protein layers Checkley et al. found that growing larval White Sea Bass at elevated pCO2 caused otoliths to be 8% and 16% larger in the 1000 and 2500 µatm treatments compared to the 430 µatm controls Control

  41. Stop !

  42. IPCC-FAR

  43. In calculating the Ksp the activity of the ionic species should be used. In practice, marine chemists would measure the concentration of Ca2+ and CO32- at which precipitation occurs. The resulting solubility product would be the apparent Ksp’ or stoichiometric constant. It’s value would depend on the conditions such as temp, pressure, and Ionic strength.

  44. 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/H2CO3 equivalence point. Modern methods involve coulombic titrations to determine end point.

  45. 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.

  46. 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 HBa H+ + Ba- H2O H+ + OH-

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