Broadly speaking, freshwater may be neutral, acid, or alkaline. The neutrality of water, or its degree of acidity or alkalinity, is known as its pH value. Much mystery has been made about the pH value of water, but, in reality, there is nothing very mysterious about it, nor is the subject so complicated, as some would have us believe. The pH value can be defined as 'a number used to express the concentration of ionised hydrogen in an aqueous fluid and is thus indicative of the reaction of that fluid, that is, the neutrality or the degree of acidity or alkalinity'. According to the theory of electrolytic dissociation all liquids of which water is a constituent contain free, positively charged hydrogen (H+) ions and negatively charged hydroxyl (OH-) ions.
When the amount of these two ions present in a liquid is exactly balanced the liquid is said to be neutral. If there be an excess of hydrogen (H+) ions the liquid is acid, and conversely if the hydroxyl (OH-) ions be in excess, it is alkaline. Absolute neutrality has a pH value of 7.07 (usually taken as 7.0). The addition of acid increases the H ion concentration; consequently the pH of all acid solutions is less than 7.07. The addition of alkali increases the concentration of the OH- ions, and decreases that of the H+ ions, so that the pH of all alkaline solutions is greater than 7.07. The range of pH values extends about equally on each side of 7.07; for the complete range of pH values forms a graduated scale from about - 0.3 to 14.5.
This scale is logarithmic; meaning that each number is ten times stronger than the preceding number. For example, a pH of 2 is ten times more acidic than a pH of 3 and one hundred times more acidic than a pH of 4. It is important to monitor pH because if a fish is forced to live in a pH level outside its preferred range, its slime coat can suffer, making it susceptible to disease. Its fecundity drops and, ultimately, the gas exchange in the gill membranes will be so reduced that the fish may suffocate.
A simple test kit which exhibit characteristic colour changes at different pH values or a hand-held electronic meter can be used to test pH. Always remove a sample of water from the aquarium to measure the pH with an electronic meter. Measurement by immersing the electrode directly in the tank can be severely compromised by other undetectable electrical currents from power filters, heaters, etc. pH electrodes which are not routinely cleaned and standardised will not provide accurate readings and will be no better than, and often far worse than, a colorimetric measurement made with the cheapest liquid-reagent test kit.
pH electrodes must be routinely checked against known pH standards to insure accuracy and need to be replaced every 9 to 12 months. The popular pocket pH "pens" are disposable meter/electrode combinations which can be inaccurate, particularly if not calibrated correctly, and do not compensate for changes in temperature. Therefore, the selection of measuring devices for pH is largely a situation in which "you get what you pay for." If you are unable to recognise the inadequacies of pH meter measurements, you are better off using dye methods. Only dyes with clear-cut colour changes around the target pH should be used.
The commonly used pH indicators for freshwater testing are bromothymol blue (yellow to green to blue as the pH increases) and phenol red (yellow to orange to red with increasing pH). The colorimetric method is the least expensive but can suffer from interferences due to discoloured water samples, salinity, organic matter, and substances, which can oxidise or reduce the reagents. In water with very low alkalinity, the indicators themselves may actually alter the pH of the sample. However, for the purposes of routine aquarium testing, colorimetric indicators are more than adequate. Some scientific supply houses now sell narrow-range litmus paper, which allows for low-cost, rapid estimation of pH.
Rainbowfishes will survive reasonably well in waters with a pH between 6 and 9. If pH readings are outside this range, growth is reduced. At values below 4 or above 10, mortalities will occur. In well-buffered aquariums with alkalinity levels above 50 mg/L, the pH will be more stable. In the morning, carbon dioxide levels are high and pH is low because of respiration during the night (carbon dioxide forms a mild acid when dissolved in water). When a suitable light source is provided, algae and other aquatic plants will produce carbohydrates and oxygen from carbon dioxide and water by photosynthesis. As carbon dioxide is removed from the water, its pH increases. In aquarium systems, the pH will generally drop in relation to the fish load, biological filtration, feeding, and maintenance schedules. Therefore, acidic water in an aquarium system is biologically different from that found in nature.
In an aquarium, acids derive primarily from two sources. The first is when carbon dioxide (directly dissolved into water or released as a respiration by-product) mixes with water to form carbonic acid.
H2O + CO2 <=> H2CO3 <=> H+ + HCO3-
water + carbon dioxide <=> carbonic acid <=> hydrogen ion + bicarbonate
The other is when ammonia undergoes nitrification by bacteria.
2 NH3 + 3 O2- > 2 NO2- + 2 H+ + 2 H2O
ammonia + oxygen —> nitrite + hydrogen ion + water
If the aquarium water is not well buffered any acid that is added serves to drive down the pH. Consequently, the daily pH swings caused by photosynthesis can combine with longer-term acid accumulations and cause the pH to suddenly drop with catastrophic results for the fish. From my own experience, most rainbowfishes in captivity do not seem to be comfortably in water below pH 6, certainly not for any extended period. A pH range of 6.5 to 7.8 is suggested for maintaining rainbowfishes in captivity.
ORP (oxidation-reduction potential) is the measure of electron activity as opposed to hydrogen activity in the case of pH. An ORP measurement is made using the millivolt mode of a pH meter. By substituting a metallic electrode for a glass pH electrode, electron activity can be detected.
© Copyright Adrian R. Tappin
Updated October, 1999.
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