Alkalinity, and the pH of Seawater

Section 8.6.2, the Permian period ended with the largest mass extinction event that has yet occurred on planet Earth. As the ocean began a sustained recovery at the beginning of the Mesozoic era, opportunities likely abounded for the survivors to take over empty ecological niches through evolutionary adaptation. Prior to the advent of planktonic 

About 25% of the carbonates deposited in shallow water are eventually eroded and carried downslope by bottom and turbidity currents to become part of the shelf and pelagic sediments. Shallow-water carbonates are also notable for their mineral composition. In addition to calcite and aragonite, some shallow-water calcifiers deposit hard parts containing high percentages of magnesium. These are referred to as magnesium-rich calcites. 

These shallow-water deposits were the sole source of biogenic carbonate until the evolution and proliferation of planktonic calcifiers, namely the coccolithophorids and the foraminiferans, around 250mybp. This enabled a shift in the site of sedimentary carbonate acciunulation from the shallow waters to the deep sea. At present, about half of the sedimentary carbonates are being buried on the shelves and the other half in the deep sea and slopes. 

Distribution of warm- and cool-water shelf carbonate In the modern-day ocean. Source From Mackenzie, F. T, and A. Lerman (2006). Carbon in the Geobiosphere Earth s Outer Shett. Springer-Veriag, p. 283. 

In terms of organic carbon generation, the coccolithophorids are a minor player, representing only 6 to 8% of global marine primary production. But their detrital remains contribute disproportionately to the burial of carbon in marine sediments. This is due to near complete loss of POC via remineralization as the detrital hard and soft parts settle to the seafloor. As estimated from Broecker s Box model in Chapter 9, only about 1% of the POM that sinks out of the surfece water is buried in marine sediments. In comparison, about 20% of the biogenic PIC survives to become buried in the sediments. 

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These simple equations and ideas provide the basis for describing the carbonate system in terms of the/coj, DIG, pH, and alkalinity of seawater. We will build up a plot similar to that in Fig. 4.1 for the important acids and bases in seawater. These are listed along with their concentrations and apparent equihbrium constants in Table 4.1. It will then be demonstrated how the constraint of charge balance (called alkalinity) determines the pH of seawater. 

Ammonium is present at very low concentrations (0.03—0.5 iM) in oceanic surface waters, at higher concentrations in coastal and estuarine waters (Sharp, 1983), and at concentrations orders of magnitude higher in sediment pore waters. In seawater, NH4+ exists as the acid base pair NH4+-NH3 (ammonia) the pFC of the pair is 9.3. The methods discussed here measure the sum of NH4+, the form that dominates at the pH of seawater ( <8.3), and NH3, the volatile form that dominates under more alkaline conditions. There are many approaches to measuring NH4+, but we win focus on the two most widely used— phenol-hypochlorite and orthophtal-dialdehyde (OPA). 

We have now described the system of equations necessary for determining the pH of seawater and the distribution of carbonate species. By including the definition and numerical value of the alkalinity to the system of equations used to determine the curves in Fig. 4.2, we have constrained the location on the plot to a single pH. The equations necessaiy to determine this location are summarized in Appendix 4.1 for the progressively more complicated definitions using the three forms of the alkalinity, Ac, Ac b. and Aj. 

Both Die and Ax increase from surface waters to the deep Atlantic, Antarctic and Pacific Oceans as one follows the route of the ocean conveyor helt (Fig. 1.12). Along this transect pH changes from about 8.2 in surface waters to 7.8 in the deep Pacific Ocean, and CO3 decreases from nearly 250 geq kg to less than a third of this value, 75 geq kg. The reason for this change has to do with the ratio of the change in Aj and DIG in the waters and is discussed in the final section of this chapter. Notice that the contribution of the nutrients Si and P to the total alkalinity is only between 0 and 5 geq kg or at most 0.2% of the total alkalinity. Although Si concentrations are much greater than those of P, the two nutrients have nearly equal contributions to the alkalinity (Table 4.4) because the pK values for two phosphate reactions are closer to the pH of seawater than is the pK for sfiicate (see Table 4.1). 

Atienza et al. [657] reviewed the applications of flow injection analysis coupled to spectrophotometry in the analysis of seawater. The method is based on the differing reaction rates of the metal complexes with 1,2-diaminocycl-ohexane-N, N, N, A/Metra-acetate at 25 °C. A slight excess of EDTA is added to the sample solution, the pH is adjusted to ensure complete formation of the complexes, and a large excess of 0.3 mM to 6 mM-Pb2+ in 0.5 M sodium acetate is then added. The rate of appearance of the Pbn-EDTA complex is followed spectrophotometrically, 3 to 6 stopped-flow reactions being run in succession. Because each of the alkaline-earth-metal complexes reacts at a different rate, variations of the time-scan indicates which ions are present. 

Figure 8.35 shows the redox state and acidity of the main types of seawaters. The redox state of normal oceanic waters is almost neutral, but they are slightly alkaline in terms of pH. The redox state increases in aerated surface waters. Seawaters of euxinic basins and those rich in nutrients (eutrophic) often exhibit Eh-pH values below the sulfide-sulfate transition and below carbonate stability limits (zone of organic carbon and methane cf figure 8.21). We have already seen (section 8.10.1) that the pH of normal oceanic waters is buffered by carbonate equilibria. At the normal pH of seawater (pH = 8.2), carbonate alkalinity is 2.47 mEq per kg of solution. 

Figure 2.12. The relationship of solubility (-log lAPcaCC ) to the log of relative coating thickness (log Z), and the variation of that relationship with time for natural seawater systems. These experiments were performed by suspending different amounts of calcite in seawater in a closed system, and monitoring pH and total alkalinity with time, g calcite cm-3 seawater = 0.001 ( ), 0.010 ( ), 0.020 (O), 0.040 (A), 0.080 ( ). (After Schoonmaker,1981.)...

The interest in the carbonate system is related to attempts to understand the uptake of fossil fuel produced CO2 by the oceans. The carbonate system can be studied by measuring pH, total alkalinity (TA), total inorganic carbon (TCO2), and the fugacity of CO2 (fco )- At least two of these variables are needed (Park, 1969) to characterize the CO2 system in the oceans. Reliable stoichiometric constants (K ) for the carbonate system are needed to determine the concentration, mol (kg solution) of the components of the CO2 system ([HCOa"], [CO2], [COa ]) and the saturation state of CaCOa as a function of salinity, temperature, and pressure (Culberson and Pytkowicz, 1968 Ingle, 1975 Millero, 1995, 2001). This includes constants for the solubility of CO2 in seawater (Weiss, 1974)... 

In order to solve the equations and determine pH and the concentrations of the species that make up the alkalinity, the apparent equilibrium constants, F, must be accurately known. These constants have been evaluated and re-evaluated in seawater over the past 50 y. The pH scales and methods of measuring pH during these experiments have been different, and this has complicated comparisons of the data until recently, when many have been converted to a common scale. Equations for the best fit to carbonate system equilibrium constants as a function of temperature and salinity are presented by Luecker et al. (2000), DoE (1994) and Millero (1995) (see Appendix 4.2). 

In continental waters, bicarbonate (I ICO ) and carbonate (CO2-) ions are the most important components of alkalinity, although in seawater other ions also contribute to alkalinity. The relative importance of HCOf and CO2- depends on the pH of the solution and can be calculated from the known dissociation constants (see Box 4.5) of these ions and the solution pH. 

The distribution of the carbon-containing species (carbonic acid, bicarbonate and carbonate) depends upon pH, as shown in Fig. 3.28 (see Box 3.4). Seawater is slightly alkaline, and its pH rarely falls outside the range 7.5-8.5 owing to the buffering effect of the equilibrium between bicarbonate and carbonate ions. Consequently, seawater contains negligible carbonic acid, and OH ions are more abundant than H+, so the main equilibria can be represented by Eqns 3.9a and 3.9b, which are combined to give Eqn 3.9c ... 

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