A Citizen's Guide to Understanding and Monitoring Lakes and Streams

Chapter 2 - Lakes

pH in Lakes

Why Is It Important?

The pH of a sample of water is a measure of the concentration of hydrogen ions. 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 concentration of the hydrogen ion. For example, there are 10 times as many hydrogen ions available at a pH of 7 than a pH of 8. The pH scale ranges from 0 to 14. A pH of 7 is considered to be neutral. Substances with pH less than 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 (lead, copper, cadmium, etc.). For example, in addition to affecting how much and what form of phosphorus is most abundant in the water, pH also determines 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 hydrogen molecules, which causes the concentration of hydrogen ions to decrease and therefore the pH to increase. 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.

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 water hovers between 6.5 and 8.5.

Expected Impact of Pollution

When pollution results in higher productivity (e.g., from increased temperature or excess nutrients), pH levels 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. There is no numerical Washington State water quality standard for pH in lakes. The standard reads there will be "no measurable change from natural conditions." (A pH of 5-6 or lower has been found to be directly toxic to fish, according to the EPA.

Generally, during the summer months in the upper portion of a productive or euthrophic 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 next section discusses secchi disk depth in lakes.

Chapter Four provides information about how to measure pH in lakes.

Temperature | Oxygen | pH | Secchi | Nutrients | Turbidity | Chlorophyll | Fecal Coliforms

Return to Table of Contents