Читать книгу Fundamentals of Aquatic Veterinary Medicine - Группа авторов - Страница 30

The term “pH” is a mathematical transformation of the hydrogen ion (H⁺) concentration; it conveniently expresses the acidity or basicity of water. The lowercase letter “p” refers to “power” or exponent, and pH is defined as the negative logarithm of the hydrogen ion concentration. Each change of one pH unit represents a 10‐fold change in hydrogen ion concentration. The pH scale is usually represented as ranging from 0 to 14, where lower number reflects higher acidity and the higher number reflects higher alkalinity. Water with a pH of 4.5 or lower has no measurable alkalinity. Water with a pH of 8.3 or higher has no measurable acidity. At 25°C, a pH value of 7 is neutral; this describes the neutral point of water at which the concentrations of hydrogen and hydroxyl ions (OH^–) are equal (each at 10–7 moles/l). The pH of freshwater ecosystems can fluctuate considerably within daily and seasonal timeframes, and most freshwater animals have evolved to tolerate a relatively wide environmental pH range. Animals can, however, become stressed or die when exposed to pH extremes or when pH changes rapidly, even if the change occurs within a pH range that is normally tolerated. The pH of water in ponds often increases during the day and decreases at night. Fish and other vertebrates have an average blood pH of 7.4. Recommended pH range for cultured fish is 6.5–9.0. Acid death point is around 4; alkaline death point is around 11. The hydrogen ion concentration affects aqueous equilibria involving ammonia, hydrogen sulfide, chlorine and dissolved metals. Interactions of pH with these variables are often more important than the direct effects of pH on aquatic animals. For example, the toxicity of ammonia to fish increases with an increase in pH. Chemical interactions among carbon dioxide, hydrogen ions and the anions that produce alkalinity buffer the pH of most natural waters in a range of about 6–8.5.