pH Information Sheet

Background

The pH of water is very important to water quality because it controls the types and rates of many chemical reactions in water, and aquatic organisms have a specific pH range in which they can live. Water (H2O) contains both hydrogen ions (H+) and hydroxyl ions (OH-). pH is a measure of the concentration of free hydrogen ions, which will indicate whether a solution is acidic or basic. Specifically, pH is equal to the negative log of the hydrogen ion concentration (or, pH = -log10[H+]). The numerical value does not have a unit (like mg/L) per se, but must be listed alongside the term pH. The values for pH are arranged on a scale from 0 to 14. A pH of 7 indicates the solution is neutral and the concentration of H+ is equal to the concentration of OH-. Values of pH less than 7 are considered acidic (more H+ are present, less OH-). Values of pH greater than 7 are considered basic (less H+ are present, more OH-). Because pH is determined based upon a log scale, each unit change in pH indicates a tenfold difference in the concentration of hydrogen ions. For example, water at pH 5 is ten times more concentrated with H+ than water at pH 6.

Natural, uncontaminated rain water is generally somewhat acidic, with a pH of about 5.6. This acidity is due to the natural dissolving of carbon dioxide (CO2) in precipitation (H2O) to form carbonic acid (H2CO3). The extra hydrogen ions are produced when the carbonic acid dissociates (breaks apart) producing H+ and bicarbonate HCO3-.

Once precipitation hits the ground, a variety of organic and inorganic chemical reactions may take place to alter the pH of water. In the upper parts of the soil, infiltrating water commonly reacts with organic matter to form organic acids, and eventually lower the value of pH (more acidic). Reaction with inorganic minerals (in rocks for example) dominate once water infiltrates beneath the soil; most of these reactions will use free hydrogen ions (buffering the solution) and therefore cause an increase in pH (more basic). The geology of a region exerts a strong control on the pH of natural waters. For example, minerals such as calcite (calcium carbonate - CaCO3), the main component of limestone and the cement that holds sandstone particles together, are especially effective at causing increases in pH. As calcium carbonate dissolves, free hydrogen ions are used. This ability to buffer, or resist the changes in pH, is called alkalinity. See the alkalinity information sheet for more background.

Once water enters lakes and streams, aquatic life may affect pH. Respiration by plants and animals and decomposition produce CO2, allowing it to react with water to form carbonic acid (eventually dissociates producing an H+) and the pH levels of a waterway can decrease (more H+). However, during daylight hours, plants photosynthesize using CO2 and keeping it from forming carbonic acid and extra H+. Under normal stream conditions, pH levels are usually highest at the end of a day of photosynthesis, lowest after a night of plant respiration.

All aquatic life has a specific pH range that it can tolerate and to which it is adapted. If the pH changes even slightly, it will stress the creatures and may even kill them. At extremely high (9.6) or low (5.0) pH values, the water becomes unsuitable for most organisms. Immature stages of aquatic insects and young fish are extremely sensitive to a pH below 5. Low pH causes an imbalance in the sodium and chloride ions in aquatic animals' blood. At low pH, hydrogen ions may be taken into cells while expelling sodium ions. Higher acidity can increase the concentration of toxic metal concentrations in a stream, such as aluminum (Al+3) and copper (Cu+2). These metals were locked up in mineral matter under neutral pH levels, but become mobile when the pH lowers. Metal can clog fish gills causing breathing complications or cause deformities to young fish. They can also settle on the stream bottom filling in spaces between rocks where insects live or eggs are laid, even smothering these eggs.

Human Impact

Acidic waters have been and continue to be a major environmental concern. Whereas unpolluted precipitation has a pH of about 5.6, the precipitation in most of the Northeast United States has a pH of between 4 and 4.5. Air pollution is the cause. Increased amounts of nitrogen oxides (NOx) and sulfur dioxides (SO2-) gases, primarily from the burning of fossil fuels by power plants and industry and from car exhaust, react with water and are converted to nitric acid (HNO3) and sulfuric acid (H2SO4) in the atmosphere. Both of these acids can dissociate to produce extra H+ and lowering the pH of the rain, and the streams that it falls and drains into. Waterways may not be affected by this acidic rain if the geology in the watershed contains a lot of acid-neutralizing rocks containing CaCO3. This type of region has high alkalinity (ability to resist changes in pH). However, a region void of this type of rock and low alkalinity can have streams that are damaged by acid rain. For example, the Adirondack region in New York has rocks and streams that are unable to neutralize the acid rain; as a result, widespread fish kills have occurred.

Coal mining operations (current and abandoned) can add acidity to a waterway through acid mine drainage (AMD). The waste material of coal mining is called spoils or overburden and is the discarded soil and crushed rock found above and between coal seams. This waste contains iron pyrite (fool's gold) and when exposed to air and water, it reacts to form iron hydroxide (Fe(OH)3) and sulfuric acid (H2SO4). The acid can dissolve other minerals and metals, and the water can become very acidic (as low as a pH of 2 even) as it enters local streams.

Water Quality Criteria

As in many chemicals, there is no distinct dividing line between safe and harmful pH levels. The drinking water standards set by the Environmental Protection Agency (EPA) calls for a minimum pH of 6.5 and a maximum pH of 8.5. Natural waters should have a pH between 5.0 and 8.5, since lower or higher values are likely to be harmful to fish populations and other aquatic life.

Example pH Data - French Creek

Data for French Creek pH concentrations were collected by the Pennsylvania Department of Environmental Protection (DEP) for a Meadville site from 1973 to present date. The pH concentrations from Meadville were tabulated and graphed for this time period and the average, maximum and minimum values were determined.

Average: pH 7.40

Maximum: pH 8.80

Minimum: 6.16

The pH trends for a typical year in French Creek shows values between 7 and 8, a range excellent for aquatic life. The pH is relatively low in the late winter through June. This is because there more acid precipitation and snow melt directly enter the stream before becoming groundwater that is buffered by bedrock and soils. Also in the spring, plants are beginning to grow, respiring more and producing CO2 that can form carbonic acid. Beginning in June, the pH begins to rise until October (drier season) because the main source of water for the stream is from well buffered groundwater. After October, another decline occurs as the stream relies less on groundwater flow and plants begin to decompose (producing CO2). Sources of acidity that enters French Creek is mainly acid precipitation (often around pH of 4.5) and organic acids (carbonic acid). However, the French Creek Watershed has the geology (CaCO3) to well buffer acid inputs, thus keeping the stream pH in the basic range. Luckily the watershed has only thin coal seams, no coal mining, and no resulting acid mine drainage.