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  pH
 

pH and Bicarbonate Buffering
The pH of a sample of water is a measure of the concentration of hydrogen ions. The term pH is derived from the French puissance d’hydrogene, meaning “strength of hydrogen,” because the hydrogen ion controls acidity (Horne and Goldman, 1994). It is calculated on a exponential scale of 1 to 14. Mathematically, pH is the negative logarithm of the hydrogen ion (H+) concentration. What this means to those of us who are not mathematicians is that at higher pH, there are fewer free hydrogen ions, and that a change of one pH unit reflects a tenfold change in the concentration of the hydrogen ion. For example, there are 10 times as many hydrogen ions available at a pH of 7 than at a pH of 8. The pH scale ranges from 0 to 14. A pH of 7 is considered to be neutral. Substances with pH of less that 7 are acidic; substances with pH greater than 7 are basic.

The pH of water determines the solubility (amount that can be dissolved in the water) and biological availability (amount that can be utilized by aquatic life) of chemical constituents such as nutrients (phosphorus, nitrogen, and carbon) and heavy metals (e.g., lead, copper, cadmium). For example, in addition to affecting how much and what form of phosphorus is most abundant in the water, pH may also determine whether aquatic organisms can make use of the phosphorus. In the case of heavy metals, the degree to which they are soluble determines their toxicity. Metals tend to be more soluble at lower pH and thus more toxic.

Reasons for Natural Variation

Photosynthesis uses up dissolved carbon dioxide. CO2 removal, in effect, reduces the acidity of the water and so pH increases. In contrast, respiration of organic matter produces CO2, which dissolves in water as carbonic acid, thereby lowering the pH. For this reason, pH may be higher during daylight hours and during the growing season, when photosynthesis is at a maximum. Respiration and decomposition processes lower pH. Like dissolved oxygen concentrations, pH may change with depth in a lake, due again to changes in photosynthesis and other chemical reactions. There is typically a seasonal decrease in pH in the lower layers of a stratified lake because CO2 accumulates. There is no light for plants to fix CO2 and decomposition releases CO2.

River water is complex; it is full of chemical "shock absorbers" that prevent major changes in pH. Small or localized changes in pH are quickly modified by various chemical reactions, so little or no change may be measured. This ability to resist change in pH is called buffering capacity. Not only does the buffering capacity control would-be localized changes in pH, but it controls the overall range of pH change under natural conditions. The pH scale may go from 0 to 14, but the pH of natural waters hovers between 6.5 and 8.5.

When carbon dioxide dissolves in water a small fraction is hydrated to form carbonic acid.

CO2 + H2O ‡ H2CO3 ‡ H+ + CaHCO3-

Because it is acidic, carbonic acid dissolves calcium carbonate rocks, neutralizing the soil and stream water, and forming calcium bicarbonate. As a result the calcium bicarbonate content of streams largely determines its pH (balance between hydrogen and hydroxide ions). Buffering is then due to the presence of carbon dioxide, carbonic acid, bicarbonate ions, and carbonate ions in freshwater and effectively promotes resistance to changes in pH.

If acid (hydrogen ions) is added to this buffer solution (water with dissolved calcium bicarbonate), the equilibrium is shifted and carbonate ions combine with the hydrogen ions to form bicarbonate. Subsequently, the bicarbonate then combines with hydrogen ions to form carbonic acid, which can dissociate into carbon dioxide and water. Thus the system pH is unaltered (buffered) even though acid was introduced.

Natural waters vary in acidity and alkalinity from local geology and other natural causes, as well as, anthropogenic inputs. Normally, pH values below 5 and exceeding 9 are harmful to organisms, thus making buffering capacity in water of critical importance.

Influence on Biota
The influence of pH and buffering on organisms is often difficult to pin down specifically due to the interrelatedness of many factors. Generally, organisms appear to thrive in waters of increasing hardness (calcium and magnesium concentration). Surveys of species abundance and species richness have shown that stream water of very low ionic concentrations has reduced species abundance and species richness. This appears to true for trout (McFadden and Cooper, 1962) in terms of growth rate and upper size limit. Molluscs, crustaceans, and leeches also show responses to varying ionic concentrations (Hynes, 1970; Macan, 1974). Russel-Hunter et al. (1967) showed that Ca2+ hardness could be associated with estimates of species richness among mollusks for geographic regions.

Ca2+ (mg/l) % of Potential Molluscan Species
3 mg/l 5%
<10 mg/l 40%
10-25 mg/l 55%
> 25 mg/l up to 100%

Often clear trends are not present when relating species abundance or species richness to ion concentrations in streams. Factors such as substrate complexity and stream temperature can interact with water chemistry and, in many cases, may be more important in determining the suitability of a stream´s environment to a particular species or taxonomic group.

 


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date last updated: Wednesday March 03 2004