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 pH: Measuring the Acidity and Alkalinity of Lakes
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Why is pH Important? 

     pH is a measure of the acidity of alkalinity of water, expressed in terms of its concentration of hydrogen ions. 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 chart below shows the pH of some common substances:

The term pH was derived from the manner in which the hydrogen ion concentration is calculated - it 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 concentrations 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 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 (lead, copper, cadmium, etc.). 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 life can use it. In the case of heavy metals, the degree to which they are soluble determines their toxicity. Metals tend to be more toxic at lower pH because they are more soluble.

Reasons for Natural Variation 

     Photosynthesis uses up dissolved carbon dioxide, which acts like carbonic acid (H2CO3) in water. 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. 

     Fortunately, lake 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, 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. 

  Expected Impact of Pollution 

     When pollution results in higher algal and plant growth (e.g., from increased temperature or excess nutrients), pH levels may increase, as allowed by the buffering capacity of the lake. Although these small changes in pH are not likely to have a direct impact on aquatic life, they greatly influence the availability and solubility of all chemical forms in the lake and may aggravate nutrient problems. For example, a change in pH may increase the solubility of phosphorus, making it more available for plant growth and resulting in a greater long-term demand for dissolved oxygen. 

 Values for pH are reported in standard pH units, usually to one or two decimal places depending upon the accuracy of the equipment used.

 Since pH represents the negative logarithm of a number,  it is not mathematically correct to calculate simple averages  or other summary statistics. Instead, pH should be reported as a median and range of values; alternatively the values could be converted to hydrogen ion concentrations, averaged, and re-converted to pH values.  Generally, during the summer months in the upper portion of a productive or eutrophic lake, pH will range between 7.5 and 8.5. In the bottom of the lake or in less productive lakes, pH will be lower, 6.5 to 7.5, perhaps. This is a very general statement to provide an example of the differences you might measure.

The Case of Acid Rain

      An important exception to the buffering of pH changes in lakes is the case of lakes affected by acid rain. Lakes that have received too much rain with a low pH (acid rain), lose their buffering capacity. At a certain point, it takes only a small bit of rain or snowmelt runoff for the pH to change. After that point, change occurs relatively quickly. According to the EPA, a pH of 5-6 or lower has been found to be directly toxic to fish (for additional information, see our acid rain links).   

 REFERENCES

      Michaud, J.P. 1991. A citizen's guide to understanding and monitoring lakes and streams. Publ. #94-149. Washington State Dept. of Ecology, Publications Office, Olympia, WA, USA (360) 407-7472.

     Moore, M.L. 1989. NALMS management guide for lakes and reservoirs. North American Lake Management Society, P.O. Box 5443, Madison, WI, 53705-5443, USA (http://www.nalms.org).

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