Magnesium concentration ocean

AH metals come originally from natural deposits present in the earth s cmst. These ore deposits result from a geological concentration process, and consist mainly of metallic oxides and sulfides from which metals can be extracted. Seawater and brines are another natural source of metals, eg, magnesium (see Chemicals frombrine Magnesium and magnesium alloys Ocean raw materials). Metal extracted from a natural source is called primary metal. 

There is some debate about what controls the magnesium concentration in seawater. The main input is rivers. The main removal is by hydrothermal processes (the concentration of Mg in hot vent solutions is essentially zero). First, calculate the residence time of water in the ocean due to (1) river input and (2) hydro-thermal circulation. Second, calculate the residence time of magnesium in seawater with respect to these two processes. Third, draw a sketch to show this box model calculation schematically. You can assume that uncertainties in river input and hydrothermal circulation are 5% and 10%, respectively. What does this tell you about controls on the magnesium concentration Do these calculations support the input/removal balance proposed above Do any questions come to mind Volume of ocean = 1.4 x 10 L River input = 3.2 x lO L/yr Hydrothermal circulation = 1.0 x 10 L/yr Mg concentration in river water = 1.7 X 10 M Mg concentration in seawater = 0.053 M. 

Magnesium concentrations as a function of depth (meters below the sea floor) in sediment porewaters from the western flank of the Juan de Fuca Ridge near 48° N in the North Pacific Ocean. The concentration decreases with depth because it is removed from solution by reaction with crustal rocks at the sediment-crustal boundary. The curves are convex upward because of porewater upwelling along the upward-flowing limb of a convection cell. Velocities of the upwelling are determined by using a one-dimensional advection-diffusion model and are indicated by the numbers on the curves. Redrafted from Wheat and MottI (2000). 

The relative standard deviation for measurements of the total alkaline earths in ocean waters is reported to be less than 0.1 % (Pate and Robinson, 1961 Culkin and Cox, 1966). In the case where an automatic titration system is applied, the author found a comparable coefficient of variation (c.v.) of 0.15 %. However, for water of lower salinity with an average magnesium content of 0.205 g/kg, a c.v. of only 0.65 % was measured by the author. Although the determination of the magnesium concentration is based on the difference between titrations, errors caused by the measurements of calcium and strontium can be neglected, when the analytical procedure described in Section 11.2.1 is used. 

Handschuh and Orgel (1973) studied the mineral struvite. It can be precipitated from ocean water in the presence of phosphate if the concentration of NH ions in the water is greater than 0.01 M. If struvite is heated with urea, magnesium pyrophosphate is obtained in a yield of about 20% after 10 days at 338 K if nucleosides are added to the reaction mixture described above, nucleoside diphosphates such as uridine-5 -diphosphate and diuridine-5 -diphosphate are formed in good yields. 

Most cation exchange occurs in estuaries and the coastal ocean due to the large difference in cation concentrations between river and seawater. As riverborne clay minerals enter seawater, exchangeable potassium and calcium are displaced by sodium and magnesium because the Na /K and Mg /Ca ratios are higher in seawater than in river water. Trace metals are similarly displaced. 

More Metals. The apparent simplicity of the problem is misleading because although the concentration of transition metal ions is small, the ocean assuredly contains trace quantities of all naturally occurring metals. We now recognize two results of coordination the properties of the metal are altered, and, equally important, the properties of the ligand are altered (coordinated ammonia is less basic, cyanide ion is less toxic) (5). Most of the catalytic activity of coordination entities recently summarized involved coordination entities of transition metal ions examples involving magnesium ion constitute the main exceptions (24). 

Seawater is unfit for drinking or agriculture because each kilogram contains about 35 g of dissolved salts. The most abundant salt in seawater is sodium chloride, but more than 60 different elements are present in small amounts. Table 14.3 lists the ions that account for more than 99% of the mass of the dissolved salts. Although the oceans represent an almost unlimited source of chemicals, ion concentrations are so low that recovery costs are high. Only three substances are obtained from seawater commercially sodium chloride, magnesium, and bromine. 

Inputs and outputs of calcium, magnesium, and carbon can be balanced for the modern ocean. This balance necessitates reverse weathering reactions in the ocean system in which CO2 consumed in weathering on land is released to the ocean-atmosphere system during the formation of minerals in the ocean. If this were not the case, regardless of juvenile emissions of CO2 to the ocean-atmosphere system, atmospheric CO2 concentrations could be reduced to vanishingly small values in less than 5000 years. 

The concentration of the water of certain mineral springs and of the ocean also affords a means of isolating salt. Less soluble constituents, such as calcium sulphate, separate first. Admixture of the salt with more soluble compounds, such as magnesium chloride, is obviated by not carrying the concentration too far. Shipper4 states that the elimination of potassium chloride can be effected by repeated crystallization of the salt from water. 

Highly saline environments are not only directly associated with present seas and oceans, but also with former seas which have led to salt deposition. These are generally hypersaline environments and may include salt lakes such as the Dead Sea, where salt concentrations may reach 4-5 M NaCl (Buchalo et al., 1998), together with salt pans and flats. In many cases, these are dominated by other ions such as potassium, magnesium, calcium, sulphate, carbonate and bicarbonate, as well as sodium and chloride. Flowers et al. (1986) estimated that about 10% of global land area was occupied by soils too saline for the growth of non-halophiles. 

The world s oceans hold 1.37x10 of water (97.2% of the total amount of water of the hydrosphere). They cover 71% of the earth s surface, are actually the biggest reservoir on our planet, and contain many important minerals. The overall content of mineral matter in the oceans is estimated to be about 5 x 10 tons [1,2]. The seas contain virtually all of the naturally occurring elements and are the only universal source of mineral wealth that is available to most nations. For some of them it is the only source. Yet, most of the elements, the microelements, are available in very low concentrations, i.e., in parts per billion (ppb). The products being extracted from seawater with economic profit at present are sodium chloride, magnesium compounds, and bromine [2-4]. During the last two decades there has been growing interest in the possibility of commercial recovery of additional minerals from seawater [5] and brines [6]. 

Several mineral phases are produced by phytoplankton and exported to depth in the ocean, associated with and analogous to the production and export of organic matter. First among these is CaCOs, comprised of two mineral phases calcite and the more soluble aragonite (Milliman, 1974 Mucci, 1983). The solubility of calcite depends also on the concentration of magnesium higher... 

The numerator of the right side is the product of measured total concentrations of calcium and carbonate in the water—the ion concentration product (ICP). If n = 1 then the system is in equilibrium and should be stable. If O > 1, the waters are supersaturated, and the laws of thermodynamics would predict that the mineral should precipitate removing ions from solution until n returned to one. If O < 1, the waters are undersaturated and the solid CaCOa should dissolve until the solution concentrations increase to the point where 0=1. In practice it has been observed that CaCOa precipitation from supersaturated waters is rare probably because of the presence of the high concentrations of magnesium in seawater blocks nucleation sites on the surface of the mineral (e.g., Morse and Arvidson, 2002). Supersaturated conditions thus tend to persist. Dissolution of CaCOa, however, does occur when O < 1 and the rate is readily measurable in laboratory experiments and inferred from pore-water studies of marine sediments. Since calcium concentrations are nearly conservative in the ocean, varying by only a few percent, it is the apparent solubility product, and the carbonate ion concentration that largely determine the saturation state of the carbonate minerals. 

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