Calcium bicarbonate

Rainwater for instance will pick up atmospheric COg and react with calcium carbonate (limestone) to form a soluble substance, calcium bicarbonate. This reaction gives water its natural hardness . 

Water impurities include dissolved and suspended soHds. Calcium bicarbonate is a soluble salt. A solution of calcium bicarbonate is clear, because the calcium and bicarbonate are present as atomic-size ions that are not large enough to reflect light. Suspended soflds are substances that are not completely soluble in water and are present as particles. These particles usually impart a visible turbidity to the water. 

The calcium carbonate precipitate was removed by filtration, and the filtered solution was found to contain 1,436 g of fructose as determined by optical rotation. A small amount of calcium bicarbonate was present as an impurity in solution and was removed by the addition of oxalic acid solution until a test for both calcium and oxalic acid was negative. The insoluble calcium oxalate precipitate was removed by filtration. 

Corrosion—various salts have different acidities (pH of brine can be controlled with lime, caustic soda, or calcium bicarbonate). 

Carbon dioxide and calcium carbonate The effect of carbon dioxide is closely linked with the bicarbonate content. Normal carbonates are rarely found in natural waters but sodium bicarbonate is found in some underground supplies. Calcium bicarbonate is the most important, but magnesium bicarbonate may be present in smaller quantities in general, it may be regarded as having properties similar to those of the calcium compound except that on decomposition by heat it deposits magnesium hydroxide whereas calcium bicarbonate precipitates the carbonate. 

Calcium bicarbonate requires excess carbon dioxide in solution to stabilise it the necessary concentration depends on the other constituents of the water and the temperature. 

The amount required to keep the calcium bicarbonate in solution. 

The mathematical relationship between carbon dioxide, calcium bicarbonate and calcium carbonate has been studied by several workers, including Langelier . The simpler form of his equation is... 

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 most common source is the supersaturation and subsequent scaling of minerals originating in the MU water. Insoluble calcium carbonate in the form of calcite (CaC03) resulting from the thermal decomposition of soluble calcium bicarbonate [Ca(HC03)2] is a classic example. Calcium carbonate quickly forms a white, friable deposit. In addition, the hydrolysis of excess bicarbonate increases... 

It should be noted that calcium bicarbonate does not exist in the solid state rather, it exists as an unstable salt in water, provided that an excess of free carbon dioxide is available to maintain equilibrium. The reaction is shown below. 

The principal temporary hardness salt in raw water is calcium bicarbonate, formed by dissolution of limestone (calcium carbonate) by... 

Partial addition of lime converts the temporary hardness salt (calcium bicarbonate) to carbonate. Fifteen to 20 ppm of the calcium carbonate remains in solution, while the remainder precipitates as an insoluble salt. Further amounts of lime convert other temporary hardness bicarbonate salts (Mg, Na, etc.) to soluble carbonates ... 

Flocculation or clarification processes are solids-liquid separation techniques used to remove suspended solids and colloidal particles such as clays and organic debris from water, leaving it clear and bright. Certain chemicals used (such as alums) also exhibit partial dealkaliz-ing properties, which can be important given that the principal alkaline impurity removed is calcium bicarbonate—the major contributory cause of boiler and heat exchanger scales (present in scales as carbonate), although closely followed by phosphate. 

Aluminum sulfate, A12(S04)3 H20, is the commonest alum used. Hydration is typically 14 to 16 H20. It hydrolyzes and polymerizes in water and typically is used within a narrow window of pH levels of 5.5 to 6.5 to minimize the solubility of aluminum in the treated water. If alkalinity is present (say, due to calcium bicarbonate), the following reaction occurs, producing insoluble aluminum hydroxide [Al(OH)3]. 

The natural supply source of carbon dioxide in MU water is primarily calcium bicarbonate alkalinity [Ca(HC03)2], which reacts under conditions of heat to form insoluble calcium carbonate and carbon dioxide. Because the precipitated carbonate cannot decompose further, no additional carbon dioxide is released. As a result, the total amount of... 

Determining calcium levels normally does not identify hardness breakthrough because the calcium salt simply reacts with phosphate precipitant (or similar treatment) and is lost as a sludge. It does, however, produce an immediate and noticeable reduction in alkalinity. (Calcium bicarbonate breaks down to calcium carbonate and carbonic acid.)... 

The carbonic acid thus formed is rich in oxygen-16. The mildly acid ground-water as well as the water of rivers and lakes, which is, therefore, also enriched in oxygen-16, dissolves limestone from surrounding rocks, to form calcium bicarbonate, which is soluble in water ... 

Whenever natural water is heated and evaporates, some of the calcium bicarbonate reconverts to calcium carbonate and reprecipitates as limestone ... 

In the ground and when in contact with rocks, the acid solution reacts with limestone, and the result is a solution of calcium bicarbonate ... 

If the water then evaporates, the dry calcium bicarbonate decomposes, recreating calcium carbonate, which precipitates and forms hard deposits and incrustations while carbon dioxide and water are released into the atmosphere ... 

The action of carbonic acid on limestone produces a calcium bicarbonate solution that is exceedingly soluble in water. (For comparison, at 20°C the solubility of calcium carbonate in water is only 0.0145 g per liter while the solubility of calcium bicarbonate is 166 g per literJ ) Magnesium ions from dolomite are also released into aqueous solution according to the same mechanism. The weathering of gypsum, calcium sulfate, also releases calcium ions into natural water supplies. 

To complete the example numerically, we take as our system a solution of 10-3 moles calcium bicarbonate dissolved in one kg (55.5 moles) of water. For... 

This compound is calcium bicarbonate or calcium hydrogen carbonate. The S042" ion is the sulfate ion. The cation is Fe2+, iron(II). 

The use of barium sulphide as a secondary sulphidizer [4] was examined on oxidized lead ores from Sicily (BaS). The results obtained were encouraging. Sulphidization using Na2S can also be improved with the use of ammonium salts (chloride and sulphate). These reagents are used in cases where the ore contains clay minerals and calcium carbonate, which prevents suphidization due to the production of soluble calcium bicarbonate. The ammonium increases the solubility of calcium carbonate and improves sulphidization. 

A low pH indicates that lime is required, but not the quantity. Excess is not only wasteful, but may render certain elements (e.g. Fe, Mn, and B) unavailable to plants. There are several methods for determining the lime requirement, including adding excess calcium bicarbonate and back-titrating the excess adding increasing amounts of calcium hydroxide and monitoring the pH and the use of a buffer solution (MAFF/ADAS, 1986, pp. 150-151), which will be described below. 

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