Alkalinity
Alkalinity refers to the capacity of a solution to neutralize acids by absorbing hydrogen ions, playing a critical role in maintaining a balanced pH level. It is quantified on a pH scale, where values above 7 indicate basic (alkaline) substances, while values below 7 indicate acidic substances. Alkalinity is influenced by various dissolved substances, including ammonia, borate, hydroxide, and carbonates, which help stabilize ecosystems by buffering against pH changes. This is particularly significant in the context of climate change, as acid precipitation—resulting from pollutants like sulfur dioxide—can lower the pH of rainwater, affecting environments that lack natural alkalinity. Such scenarios often occur in regions with granite-based soils, which are less capable of neutralizing acidic inputs. Maintaining adequate alkalinity in bodies of water is vital for protecting aquatic life, as it mitigates the impacts of acid rain and supports ecosystem health. To enhance alkalinity, substances such as limestone are occasionally added to water bodies, ensuring a sustainable environment amid changing climate conditions.
Subject Terms
Alkalinity
Definition
Alkalinity is a measurement of the capacity of a solution to neutralize acid by taking up hydrogen atoms. On the pH scale, which rates the alkalinity or acidity of a given substance, alkalines, or bases, are denoted by numbers greater than 7, whereas acids are denoted by numbers less than 7. Some substances that contribute to a solution’s alkalinity are dissolved ammonia, borate, hydroxide, nitrate, phosphate, silicate, and sulfide. The alkalinity of a substance can help protect the balance of the substance, as naturally basic substances can effectively neutralize or recover from the addition of acids.
Significance for Climate Change
Rainwater is normally a weak carbonic acid solution. In the atmosphere, the water molecules in rain mix with molecules, which, because of their weak bonds, can then form hydrogen and bicarbonate ions. Acid precipitation results when pollutants, such as sulfur dioxide or nitrogen oxides, remove low-pH acids from the atmosphere in the form of rain, snow, sleet, or hail. If water or soil where this acidic precipitation falls lacks natural alkalinity, as is the case with soils based on granite or other hard rocks with low carbonate content, the water or soil will be unable to neutralize the acid and the pH balance of the water or soil may be affected, altering the dynamics of the ecosystem.
The alkalinity of a body of water or soil can act as a buffer that can prevent drastic changes in the pH balance and, thus, can more easily recover from the addition of any type of acid. Thus, mildly basic bodies of water protect aquatic life and are less vulnerable to acid rain. The carbon in carbonate rocks, such as limestone, acts as a hydrogen absorber. Often, to increase the alkalinity of a body of water and thus to protect against fluctuations in the pH balance leading to algal bloom, (also called limestone) is added.
Bibliography
Brimblecombe, P., et al., eds. Acid Rain: Deposition to Recovery. New York: Springer, 2007.
Jenkins, J. C., K. Roy, C. Driscoll, and C. Buerkett. Acid Rain in the Adirondacks: An Environmental History. Ithaca, N.Y.: Cornell University Press, 2007.
Lehmann, Nele. "Alkalinity Responses to Climate Warming Destabilize the Earth’s Thermostat." Nature Communications, vol. 14, no. 1648, 24 Mar. 2023, doi.org/10.1038/s41467-023-37165-w. Accessed 11 Dec. 2024.
Schindler, D. W. “Effects of Acid Rain on Freshwater Ecosystems.” Science 239, no. 4836 (January 8, 1988): 149-157.