Bicarbonates are one of the most common dissolved substances found naturally in water supplies across the globe. They represent a fundamental chemical component of water. The presence of these compounds influences everything from a water body’s environmental stability to its taste and how it interacts with plumbing systems. Understanding the bicarbonate ion (\(\text{HCO}_3^-\)) and its function is important for grasping the overall quality and chemical behavior of drinking water.
The Chemistry and Source of Bicarbonates
The bicarbonate ion is a polyatomic anion that forms a part of the dissolved inorganic carbon system in water. This ion is created when carbon dioxide (\(\text{CO}_2\)) dissolves in water, initiating a series of chemical reactions. The presence of bicarbonates is a direct result of water’s journey through the earth’s crust.
Rainwater naturally absorbs carbon dioxide from the atmosphere as it falls, which creates a weak solution of carbonic acid (\(\text{H}_2\text{CO}_3\)). As this slightly acidic water percolates through the ground, it encounters carbonate-rich rocks, such as limestone (\(\text{CaCO}_3\)). The carbonic acid then reacts with the limestone, dissolving it and forming soluble calcium bicarbonate (\(\text{Ca}(\text{HCO}_3)_2\)).
This geological process is the main mechanism by which bicarbonates enter water supplies. Bicarbonates are highly soluble, remaining dissolved in the water as it moves through aquifers and waterways. This natural weathering and dissolution process links the carbon cycle with the hydrologic cycle, making bicarbonates a ubiquitous component of freshwater chemistry.
Bicarbonates as the Key to Water Alkalinity
Bicarbonates are the primary contributors to water alkalinity, which is a measure of the water’s capacity to neutralize acids. Alkalinity should not be confused with \(\text{pH}\), which measures the concentration of hydrogen ions. Alkalinity, instead, is a measure of the water’s reserve capacity to resist changes in \(\text{pH}\) when an acid is introduced.
Bicarbonates, alongside carbonate and hydroxide ions, act as a natural buffer system within the water. The buffering action is the ability of a solution to absorb excess hydrogen ions (\(\text{H}^+\)) or donate them, thereby stabilizing the \(\text{pH}\). In natural waters, bicarbonates achieve this by readily reacting with any incoming acid.
When an acidic substance is added, the bicarbonate ions absorb the excess hydrogen ions, forming carbonic acid. This consumption of hydrogen ions prevents the water’s \(\text{pH}\) from dropping significantly. Conversely, if a base is added, the carbonic acid can release hydrogen ions back into the water, preventing the \(\text{pH}\) from rising too high.
This buffering system is important for aquatic life and environmental stability. A stable \(\text{pH}\) environment protects organisms from sudden and potentially harmful shifts in acidity caused by acid rain or chemical runoff. Water bodies with low alkalinity are poorly buffered and are much more susceptible to drastic \(\text{pH}\) changes, which can be devastating to ecosystems.
Bicarbonates and the Practical Effects on Water Quality
The concentration of bicarbonates in drinking water has several practical implications for consumers, most notably concerning water hardness and taste. Bicarbonates are directly responsible for “temporary hardness,” which is caused by the presence of dissolved calcium and magnesium bicarbonates.
Temporary hardness is distinguishable because it can be removed simply by heating the water. When water containing these bicarbonates is heated, they decompose and precipitate out of the solution. This process forms insoluble calcium carbonate, commonly known as limescale, which is the hard, white deposit seen inside kettles, water heaters, and pipes.
This formation of limescale can decrease the efficiency of heating elements and eventually clog plumbing. In contrast, “permanent hardness” is caused by non-bicarbonate salts, such as calcium and magnesium sulfates or chlorides, which do not precipitate out upon boiling.
Bicarbonates also affect the sensory quality of water. The presence of these mineral compounds contributes to the pleasant and neutral taste of drinking water, often associated with high-quality mineral water. A moderate level of bicarbonates is generally preferred over the flat taste of de-mineralized water.
In water quality reports, bicarbonates are usually quantified as part of the total alkalinity measurement. Total alkalinity is measured by titration with a strong acid and is reported in units of milligrams per liter (\(\text{mg/L}\)) or parts per million (\(\text{ppm}\)) as calcium carbonate equivalent (\(\text{CaCO}_3\)). This measurement provides a standardized way to assess the water’s buffering capacity, which is dominated by its bicarbonate content.