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.
for Natural Variation
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.
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.
Impact of 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
for pH are reported in standard pH units, usually to one or two decimal
places depending upon the accuracy of the equipment used.
pH represents the negative logarithm of a number, it is not
mathematically correct to calculate simple averages or other
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.
during the summer months in the upper portion of a productive or eutrophic
lakes, 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.
Case of Acid
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
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.
M.L. 1989. NALMS management guide for lakes and reservoirs. North
American Lake Management Society, P.O. Box 5443, Madison, WI, 53705-5443,