Carbonate and Bicarbonate in Water
Carbonate and Bicarbonate in Water differ by their electrical charge. Carbonate has a CO3 -2 charge while Bicarbonate has a HCO3 -1 charge.
Carbon dioxide (CO2) can be dissolved by naturally circulating waters with air and dissolving rock, sediment and soils. Depending upon pH, CO2 appears principally as bicarbonate and carbonate ions.
Carbonate that follows this path represents a linkage between the carbon cycle and the hydrologic cycle.
The large supply of atmospheric carbon dioxide is partly intercepted by photosynthesizing vegetation. They convert it to cellulose starch and related carbohydrates. These products are later reduced via respiration to carbon dioxide and water with a release of stored energy. The concentration of carbonates in natural waters is a function of dissolved carbon dioxide, temperature, pH, cations and other dissolved salts.
At one atmosphere of pressure, pure C02 gas over distilled water will produce a solution that would have a pH near 3.6. In the presence of excess calcite, however, the solution would contain some 350 mg/L of dissolved calcium and its pH would be near 6.
A considerable part of the released carbon dioxide must return directly to the atmosphere and obviously the efficiency of utilization of the dissolved portion will be far below the theoretical maximum.
Water quality records for streams in the United States indicate that the maximum rate for calcium and bicarbonate removal is near 400 tons per square mile per year, but most streams carry far less than half this much. These figures do suggest that under most favorable conditions, limestone may be rather rapidly eroded.
It may also be of interest to note that a considerable part of the anionic load of many streams is a contribution from carbon dioxide of the atmosphere rather than from the rocks of the drainage basin.
Bicarbonate concentration of natural waters generally is held within a moderate range by the effects of the carbonate equilibria.
Most surface streams contain less than 200 mg/L Carbonate and Bicarbonate, but in ground water somewhat higher concentrations are not uncommon. Concentrations over 1,000 mg/L sometimes occur in waters which are low in calcium and magnesium and especially where processes releasing carbon dioxide such as sulfate reduction are occurring in the ground water reservoir.
Many of the carbonates are quite insoluble in water. Generally more so than the chlorides, nitrates or sulfates. There is a tendency for certain carbonate salts to be removed by precipitation or absorption.
In more calcareous environments, the circulation of water rich in carbon dioxide produce solutions that are highly supersaturated when exposed to the air. Such solutions deposit large quantities of calcium carbonate as travertine near their points of discharge.
In hard waters, particularly in headwaters that are fed by limestone springs, deposits of calcium carbonate are often layed down. These may form large solid structures that can dam up a stream or produce waterfalls.
This precipitated material is travertine. Some of the deposition is probably purely chemical and is caused by the loss of equilibrium carbon dioxide necessary to keep the calcium carbonate and bicarbonate in solution. However, it is nearly always associated with algae and to a lesser extent with mosses that cause deposition of calcium carbonate by photosynthesis.
As will be discussed in alkalinity, it is important to note that bicarbonates tend to reach equilibrium with the carbonates. For this reason, most running waters are “bicarbonate” waters in a limnological sense and show the complicated relationships between
pH, C02, H2CO3, H+, C03–, HCO3-, calcium +2 and magnesium +2.
It is sufficient to stress only a few points:
(1) Rainwater reaching the water courses as a runoff from bogs, dense forest litter, and similar substrata tends to have a low pH because of the hydrogen ions produced by disassociation of carbonic acid and the loss of cations by base exchange with the organic matter.
(2) Water which has percolated through the soil is also rich in carbon dioxide and similarly tends to be rich in hydrogen ions according to this equation:
H20 + C02 -> H2CO3 ⇔ H + HCO3
(3) Calcium carbonate which is a common constituent of many rocks is almost insoluble in water, but it dissolves fairly readily as bicarbonate in carbonic acid. It neutralizes soil water where it occurs, according to this reaction:
CACO3 + H2CO3 -> cA(Hco3>2 ⇔ CA+2 + 2HC03-2.
(4) Calcium bicarbonate in solution is a good buffer system. It resists change in pH, but it remains in solution only in the presence of a certain amount of free carbon dioxide. Any process which removes carbon dioxide, as does photosynthesis, tends to cause precipitation of calcium carbonate from solution, especially where the bicarbonate is abundant.
Therefore, springs in limestone regions are often very rich in calcium bicarbonate where they emerge to the surface. As the water flows along, it looses carbon dioxide to the atmosphere and by photosynthesis. After some distance this loss becomes the loss of equilibrium C02 and the deposition of calcium carbonate occurs by a reversal of the first part of the equation under #3 above.
This process is also aided by the ability of many plants to make direct use of bicarbonate ions as shown in this equation:
CA(HCO3)2 -> CAC03 + H20 + C02 (photosynthesized).
Deposition of calcium carbonate is therefore a common feature of streams in limestone areas and is a subject in the alkalinity section.
The carbon dioxide released within the soil by respiration.Decay is capable of producing low pH in circulating water if minerals that act as proton acceptors are scarce. Soils of humid, temperate regions may become depleted in calcium carbonate by leaching and the pH of ground water at shallow depths may be rather low.
In general, it may be expected that carbonates in themselves are not detrimental to fish life, but their buffering action and effect upon pH may contribute to the toxicity of high pH values.
Carbon dioxide may be a limiting factor in soft waters where no bicarbonate ions are available because water in equilibrium with normal air containing 4.0 ml/L contains only 0.68 ml/L in solution of carbon dioxide. This is much less than is available to land plants. This difference is one of the chief reasons for many of the peculiarities of submerged aquatic plants.
To summarize carbonate and bicarbonate, if a spring is from limestone or some other very calcareous rock, the water will be heavily charged with calcium bicarbonate. It will loose carbon dioxide very rapidly and its pH will rise. After some distance, the loss of carbon dioxide to the atmosphere and by photosynthesis will lead to the deposition of calcium carbonate. This process will decline steadily as equilibrium is attained.
In this well buffered hard water the pH will not rise above about 8.3 even at times of very active photosynthesis. These changes will occur while the water flows only a very short distance.
Water from noncalcareous springs will similarly loose carbon dioxide, increase its pH and acquire oxygen in quite a short distance. If the source of the spring is a big, acid swamp, swampy woodland, or rain forest, the water may contain ferrous bicarbonate. As pH rises and oxygen is acquired, ferric hydroxide will be deposited probably over a distance measurable only in tens or hundreds of meters.
Unless the water remains acidic, little iron will remain in solution. This part of the stream could be coated with rust colored masses of iron bacteria. If this water is soft, its pH will fluctuate markedly because of photosynthesis and it may exceed 8.3 in the daytime.
End of Carbonate and Bicarbonate in WaterMore about specific cations and anions…
Sulfate in Water
Chloride in Water
Fluoride in Water
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Primary Nutrients in Water and Eutrophication
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