Calcium carbonate equilibrium with dissolved

The most actively cycled reservoir of carbon is atmospheric C02 (it constitutes 0.034% of the atmosphere). Carbon dioxide dissolves readily in water and is in direct equilibrium with dissolved inorganic forms of carbon (H2C03, HCO, and CO7-, see Section 6.2.1.3). Once there, it may precipitate as solid calcium carbonate (limestone). Corals and algae encourage this reaction and build up limestone reefs in the process, but a much larger portion in the deep sea equilibrates only at the slow rate of... 

The most important property of the dissolved solids in fresh waters is whether or not they are such as to lead to the deposition of a protective film on the steel that will impede rusting. This is determined mainly by the amount of carbon dioxide dissolved in the water, so that the equilibrium between calcium carbonate, calcium bicarbonate and carbon dioxide, which has been studied by Tillmans and Heublein and others, is of fundamental significance. Since hard waters are more likely to deposit a protective calcareous scale than soft waters, they tend as a class to be less aggressive than these indeed, soft waters can often be rendered less corrosive by the simple expedient of treating them with lime (Section 2.3). 

The effect of pH on the corrosion of zinc has already been mentioned (p. 4.170). In the range of pH values from 5 -5 to 12, zinc is quite stable, and since most natural waters come within this range little difficulty is encountered in respect of pH. The pH does, however, affect the scale-forming properties of hard water (see Section 2.3 for a discussion of the Langelier index). If the pH is below the value at which the water is in equilibrium with calcium carbonate, the calcium carbonate will tend to dissolve rather than form a scale. The same effect is produced in the presence of considerable amounts of carbon dioxide, which also favours the dissolution of calcium carbonate. In addition, it is important to note that small amounts of metallic impurities (particularly copper) in the water can cause quite severe corrosion, and as little as 0-05 p.p.m. of copper in a domestic water system can be a source of considerable trouble with galvanised tanks and pipes. 

Figure 29.2 shows the mineralogic results of the calculation. Dolomite dissolves, since it is quite undersaturated in the waste fluid. The dissolution adds calcium, magnesium, and carbonate to solution. Calcite and brucite precipitate from these components, as observations from the wells indicated. The fluid reaches equilibrium with dolomite after about 11.6 cm3 of dolomite have dissolved per kg water. About 11 cm3 of calcite and brucite form during the reaction. Since calculation... 

Egg shells are made of calcium carbonate, CaCCE. The chicken ingeniously makes shells for its eggs by a process involving carbon dioxide dissolved in its blood, yielding carbonate ions which combine chemically with calcium ions. An equilibrium is soon established between these ions and solid chalk, according to... 

Acid Hydrolysis. The water that enters soil as rain or snow is in equilibrium with CO2 in the atmosphere, which dissolves to form carbonic acid. Unpolluted rainwater has a pH of approximately 5.7, whereas water in soil pores may be exposed to air containing a higher partial pressure of CO2 than the free atmosphere, and hence soil water may be more acidic (see Section 5.4). It is the attack on soil minerals by this weak carbonic acid that is the major chemical weathering process in most soils. For example, acid hydrolysis of calcium carbonate yields calcium and bicarbonate ions ... 

Calcium Carbonate Protective Scale. The LangeHer saturation index (LSI) is a useful tool for predicting the tendency of a water to deposit or dissolve calcium carbonate. Work pubHshed in 1936 deals with the conditions at which a water is in equilibrium with calcium carbonate. An equation developed by LangeHer makes it possible to predict the tendency of calcium carbonate either to precipitate or to dissolve under varying conditions. The equation expresses the relationship of pH, calcium, total alkalinity, dissolved soHds, and temperature as they relate to the solubiHty of calcium carbonate in waters with a pH of 6.5—9.5 ... 

Figure 26-18 shows the carbon cycle. Carbon dioxide in the atmosphere is in equilibrium with an enormous quantity that is dissolved in oceans, lakes, and streams. Some of this dissolved CO2 was once in the form of calcium carbonate (CaC03), the main component of the shells of ancient marine animals. The shells were eventually converted into limestone, which represents a large store of carbon on Earth. When the limestone was exposed to the atmosphere by receding seas, it weathered under the action of rain and surface water, producing carbon dioxide. Some of this CO2 was released into the atmosphere. This process continues today. 

An interesting paradox about seawater is that calcium carbonate does not spontaneously precipitate, but neither do sea shells on the beach dissolve. This suggests that the oceans are not far from equilibrium with respect to the system CaC03(s)/Ca-+(aq)/ C03 (aq). However, this precipitation is not confirmed by calculations unless considerable care is taken. 

Subareal plants use atmospheric (gaseous) C02 (C02(g)) as their photosynthetic carbon source, which has a mean 813C value of c— 7%o. Subaquatic plants use the dissolved (aqueous) C02 (C02(aq)), which is at one end of the series of equilibria shown in Eqn 3.8. Both these assimilatory processes are accompanied by isotopic fractionations, as discussed in Section 5.8.2. In marine environments there are also C isotopic fractionations associated with the formation of calcium carbonate tests (using bicarbonate) by some organisms that for formation of calcite is different from that for aragonite. The overall fractionation, caicjee co2(aq) s l31 6 and temperature dependent (Fig. 5.55 Mook et al. 1974 Morse Mackenzie 1990), primarily because of the equilibrium between dissolved C02 and bicarbonate... 

An alternative method is to add a small amount of Mg " to the EDTA solution. This immediately reacts with EDTA to form MgY with very little free Mg " in equilibrium. This, in effect, reduces the molarity of the EDTA. So the EDTA is standardized after adding the Mg " by titrating primary standard calcium carbonate (dissolved in HCl and pH adjusted). When the indicator is added to the calcium solution, it is pale red. But as soon as the titration is started, the indicator is complexed by the magnesiiim and turns wine red. At the end point, it changes to blue, as the indicator is displaced from the magnesium. No correction is required for the Mg added because it is accounted for in the standardization. This solution should not be used to titrate metals other than calcium. 

Reaction with carbon dioxide. The increase in solubility of limestones in the presence of carbon dioxide is due to reversible chemical reaction (3.1) and (3.2), which form calcium and magnesium bicarbonates. For example, at 20 °C approximately 30 mg/1 of calcite will dissolve in distilled water at equilibrium with the atmospheric carbon dioxide [3.9]. 

A knowledge of the position of chemical equilibria is of the utmost importance to us in water chemistry. By knowing the position of chemical equilibria, we can determine whether it is possible for certain reactions between reactants at given concentrations to proceed. For example, we can provide answers to questions such as Will calcium carbonate tend to precipitate or dissolve in this water Can I possibly oxidize sulfide with nitrate and so on. There are two general ways to answer questions like these. The first is to do an experiment and the second is to calculate the answer using previously determined equilibrium data. Although the first way may be more enjoyable to those who like puttering around in the laboratory, the second approach is far superior if time is of the essence. 

If the water in contact with cement is in equilibrium or is oversaturated with respect to calcimn carbonate, its deposition leads to blocking of the pores. The calcium carbonate deposit is then stable and the cement material is protected for as long as the transported water remains under these conditions. If the transported water has a low total inorganic carbon (TIC), the deposit m not lead to blocking of the pores, regardless of whether the water transported is oversaturated with CaCOs or not. In this case the amoimt of CaCOs precipitation possible is smaller than the amount of calcium species dissolved. 

The two parameters that control corrosivity of soft waters are the pH and the dissolved oxygen concentration. In hard waters, however, the natural deposition on the metal surface of a thin diffusion-barrier film composed largely of calcium carbonate (CaCOs) protects the underlying metal. This film retards diffusion of dissolved oxygen to cathodic areas, supplementing the natural corrosion barrier of Fe(OH)2 mentioned earlier (Section 7.2.3). In soft water, no such protective film of CaCOs can form. But hardness alone is not the only factor that determines whether a protective film is possible. Ability of CaCOs to precipitate on the metal surface also depends on total acidity or alkalinity, pH, and concentration of dissolved solids in the water. For given values of hardness, alkalinity, and total dissolved salt concentration, a value of pH, given the symbol pHs, exists at which the water is in equilibrium with solid CaCOs. When pH > pHs, the deposition of CaCOs is thermodynamically possible. 

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