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pH and Water Quality
The acidity or basic nature of a solution is expressed as the pH. The concentration of the
hydrogen ion [H+] in a solution determines the pH. Mathematically this is expressed as:
pH = - log [H+]
The pH value is the exponent to the base 10 of the hydrogen ion concentration. The more acidic
the solution, the lower the pH; the more basic, the higher the pH. Each change in pH unit
represents a tenfold change in acidity. For example, a solution at pH 3 is ten times more acidic
than one at pH 4.
The pH of a water body results from the ratio of H+ to OH- ions. In natural waters this usually is
dependent on the carbonic acid equilibrium. When carbon dioxide from the air enters freshwater,
small amounts of carbonic acid are formed which then dissociate into hydrogen ions and
bicarbonate ions, as shown in the following equations.
CO2 + H2O ===> H2CO3 (carbonic acid)
H2CO3 ===> HCO31- + H+
This increase in H+ ions makes the water more acidic and lowers the pH. If CO2 is removed (as in
photosynthesis) the reverse takes place and pH rises. This process is also related to the presence of
carbonates, calcium, or other ions such as magnesium as discussed under alkalinity. (Note: H+
ions actually occur as hydronium ions H3O + [hydrated protons]; however, for simplicity, they
have been represented as H+.)
Methodology: In the laboratory, pH is an electrometric measurement. A computer aided
titrimeter (CAT) is used to determine the pH value. This value is determined at a particular millivolt
reading on the instrument. By using indicators that change color, pH values may also be determined.
Phenolphthalein is an end-point indicator used in titration procedures, while litmus is commonly used
in paper-coded pH measurements.
In the field, a simple colorimetric method will be employed unless an electronic field pH meter is
available.
Environmental Impact: A pH range of 6.0 to 9.0 appears to provide protection for the life of
freshwater fish and bottom dwelling invertebrates.
THe table below gives some special effects of pH on fish and aquatic life.
Table III
| Limiting pH Values
| | Minimum
| Maximum
| Effects
| | 3.8
| 10.0
| Fish eggs could be hatched, but
deformed young were often produced | | 4.0
| 10.1
| Limits for the most resistant fish
species | | 4.1
| 9.5
| Range tolerated by trout | | 4.3
| ---
| Carp died in five days | | 4.5
| 9.0
| Trout eggs and larvae develop
normally | | 4.6
| 9.5
| Limits for perch | | 5.0
| ---
| Limits for stickleback fish | | 5.0
| 9.0
| Tolerable range for most fish | | ---
| 8.7
| Upper limit for good fishing waters | | 5.4
| 11.4
| Fish avoided waters beyond these
limits | | 6.0
| 7.2
| Optimum (best) range for fish eggs | | 1.0
| ---
| Mosquito larvae were destroyed at this
pH value | | 3.3
| 4.7
| Mosquito larvae lived within this range | | 7.5
| 8.4
| Best range for the growth of algae |
One of the most significant environmental impacts of pH is the affect that it has on the solubility, and
thus the bioavailability, of other substances. This process is important in surface waters. Runoff from
agricultural, domestic, and industrial areas may contain iron, lead, chromium, ammonia, mercury or
other elements. The pH of the water affects the toxicity of these substances. As the pH falls
(solution becomes more acidic) many insoluble substances become more soluble and thus available
for absorption. For example, 4 mg/L of iron would not present a toxic effect at a pH of 4.8.
However, as little as 0.9 mg/L of iron at a pH of 5.5 can cause fish to die.
Criteria: For aquatic life the pH should be between 6.0 and 9.0 pH units.
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