Seawater calcite solubility

A classic example of metastability is surface-seawater supersaturation with respect to calcite and other carbonate minerals (Morse and Mackenzie 1990 Millero and Sohl 1992). The degree of calcite supersaturation in surface seawater varies from 2.8- to 6.5-fold between 0 and 25 °C (Morse and Mackenzie 1990). In Fig. 3.18, experimental calcite solubility (metastable state) is approaching model calcite solubility (stable state) at subzero temperatures. In Table 5.1, the difference in seawater pH, assuring saturation or allowing supersaturation with respect to calcite, is 0.38 units. Moreover, in running these calculations, it was necessary to remove magnesite and dolomite from the minerals database (Table 3.1) because the latter minerals are more stable than calcite in seawater. But calcite is clearly the form that precipitates... 

In the previous chapter, the fact that stoichiometric and apparent constants have been widely used in seawater systems was discussed. Berner (1976) reviewed the problems of measuring calcite solubility in seawater, and it is these problems, in part, that have led to the use of apparent constants for calcite and aragonite. The most difficult problem is that while the solubility of pure calcite is sought in experimental seawater solutions, extensive magnesium coprecipitation can occur producing a magnesian calcite. The magnesian calcite should have a solubility different from that of pure calcite. Thus, it is not possible to measure pure calcite solubility directly in seawater. 

Schoonmaker J.E. (1981) Magnesian calcite-seawater reactions solubility and recrystallization behavior. Ph.D. dissertation, Northwestern Univ., Evanston, IL. 

Hiickel theory [or the Giintelberg or Davies equation (Table 3.3)] may be used to convert the solubility equilibrium constant given at infinite dilution or at a specified / to an operational constant, valid for the ionic strength of interest. In seawater solubility equilibrium constants, experimentally determined in seawater, may be used. For example, the CaC03 calcite solubility in seawater of specified salinity may be defined by = [Ca " ] [CO f ], where [Caj ] and [C03f ] are the total concentrations of calcium and carbonate ions, for example,... 

Example 7.9. Effect of Pressure and Temperature on Calcite Solubility in Seawater An enclosed sample of the surface seawater, as discussed in Example 7.8 (25°C pH = 8.2 [Ca ] = 1.06 x 10" M [Carb-Alk] = 2.4 x 10" eq liter" ), is cooled to 5°C and then subjected to increases in total pressure of up to 1000 atm (equivalent to exposing the sample to increased water depths of approximately 10,000 m). How does the composition, pH, [CO3"], [Ca "], and extent of oversaturation change as a result of the temperature change at 1 atm and as a result of the pressure change at 5°C The water is incipiently oversaturated with respect to calcite. Assume that CaC03... 

Schoonmaker, J. E. (1981), Magnesian Calcite-Seawater Reactions Solubility and Recrystallization Behavior, Ph.D. Dissertation, Northwestern University, Evanston, IL. 

Certain ocean precipitates incorporate the Mg2+ ion in their makeup. In calcite (CaC03), it has been shown that the magnesium content enhances calcite solubility, which, in turn, slows crystal growth. This raises concerns about increasing the magnesium content of seawater (as from oil production), which can have adverse implications for calcareous marine organisms such as plankton. 

The alkalinity and DIC data were used to calculate the carbonate ion concentration and pH of the culture medium during the second half of the first experiment, and throughout the second experiment (Fig. 3). These calculations used the program C02SYS (Lewis Wallace 1998) with the carbonate dissociation constants of Roy et al. (1993), the calcite solubility of Mucci (1983), and the assumption that the boron/salinity ratio of the culture system water was equal to the seawater ratio. Because much of the culture system water in both years was Instant Ocean, it may not be correct to estimate the total borate concentration from the whole-ocean boron/salinity relationship. However, trends in the concentrations of carbonate system species during each year will be independent of the actual absolute total borate concentration. 

It is doubtful that formation and dissolution of any mineral in low temperature aqueous solutions has been more fully investigated than the magnesian caicite. This mineral is a preponderant carbonate phase, mostly of biogenic origin, in seawater. Fig. 8.8 gives some data on the solubilities of Mg-calcites as a function of MgC03 content. 

All surface seawater is presently supersaturated with respect to biogenic calcite and aragonite with Cl ranging from 2.5 at high latitudes and 6.0 at low latitudes. The elevated supersaturations at low latitude reflect higher [COj ] due to (1) the effect of temperature on CO2 solubility and the for HCO3, and (2) density stratification. At low latitudes, enhanced stratification prevents the upwelling of C02-rich deep waters. 

As with the calcareous tests, BSi dissolution rates depend on (1) the susceptibility of a particular shell type to dissolution and (2) the degree to which a water mass is undersaturated with respect to opaline silica. Susceptibility to dissolution is related to chemical and physical factors. For example, various trace metals lower the solubility of BSi. (See Table 11.6 for the trace metal composition of siliceous shells.) From the physical perspective, denser shells sink fester. They also tend to have thicker walls and lower surface-area-to-volume ratios, all of which contribute to slower dissolution rates. As with calcivun carbonate, the degree of saturation of seawater with respect to BSi decreases with depth. The greater the thermodynamic driving force for dissolution, the fester the dissolution rate. As shown in Table 16.1, vertical and horizontal segregation of DSi does not significantly coimter the effect of pressure in increasing the saturation concentration DSi. Thus, unlike calcite, there is no deep water that is more thermodynamically favorable for BSi preservation they are all corrosive to BSi. 

The other reason why the average salinity of seawater is 35%o lies in the fundamental chemistry of major ions. For example, the sevenfold increase in the Na /K ratio between river water and seawater (Table 21.8) reflects the lower affinity of marine rocks for sodium as compared to potassium. In other words, the sodium sink is not as effective as the one for potassium. Thus, more sodium remains in seawater, with its upper limit, in theory, being controlled by the solubility of halite. Likewise, the Ca /Mg ° ratio in seawater is 12-fold lower than that of river water due to the highly effective removal of calcium through the formation of biogenic calcite. 

The central question was, how major is the influence of coprecipitated magnesium Berner (1976) used a clever approach to find out. His reasoning was that whereas calcite has a major coprecipitate from seawater (Mg) capable of altering its solubility, aragonite that can be precipitated from seawater is greater than 99% pure, and its solubility in seawater should be measurable. He also pointed out that if the solubility ratio of calcite to aragonite was precisely known, it would... 

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.)...

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