Acidification and Buffering in Ground- and Freshwaters
In the absence of strong acid anions such as SO2- and NO-, pure rainwater has a pH of 5.6-5.7. The latter is caused by the equilibrium with atmospheric CO2 (see the section titled 'Introduction').
A major part of rainwater reaching the groundwater and/or surface waters percolates through soils and bedrock. A small portion of (acidified) rainwater is directly deposited into lakes and streams. Areas highly affected by water acidification are small watersheds with shallow soil cover, rapid flushing rates, and slowly weathering bedrock, such as granite and quartzite. These types of soil and bedrock do not contain unstable or readily soluble minerals such as CaCO3 and MgCO3, which are very effective in neutralizing the acids (see the section titled 'Carbonate dissolution').
The natural buffering system in surface waters is provided by HCO-, released by the weathering of soil/ bedrock minerals, and by the balance between dissolved atmospheric CO2 and CO2 from respiration/decomposition processes (Box 2). As for soils, freshwater and groundwater acidification can be defined as a decrease in acid neutralization capacity (ANCaqua). Because of the electroneutrality constraint in solutions, ANCaqua can also be defined as the sum of all 'base' cations minus the sum of all 'strong' acid anions (Box 10).
In acidified surface waters, the pH and/or ANCaqua have fallen significantly below pre-industrial levels (ANCaqua < 0.10 mmolc of HCO- per liter). Elevated levels of Al compounds and low pH values usually accompany low ANCaqua. One visible sign of acidification is that the water becomes clearer because humus substances that normally color the water precipitate out together with Al (or Fe) compounds. Biomass production (algae and bacteria) and decomposition slow down, and organic matter such as leaves often collect on the lake or riverbeds.
Box 9 Aluminum hydroxides, iron oxides and hydroxides
Al(OH)3 4 3H4 $ Al34 4 3H2O pH range: 4.2-3.2 FeOOH 4 3H4 $ Fe34 4 2H2O pH range: <3.2
Of all CO2 emitted globally due to land-use changes, fossil fuel burning, and cement production in the past 200 years, only about half has remained in the atmosphere. Besides
Box 10 Acid neutralization capacity in freshwaters - ANCaqua
Constraint of electroneutrality:
HCO- + OH- + Cl- + NO3- + 2SO4- = H+ + Al3+ + Al(OH+) + 2Ca2+ + 2Mg2+ + K+ + Na+
HCO3- + OH- - H- - Al3+ - Al(OH+) = 2Ca2+ + 2Mg2+ + K+ + Na+ - Cl- - NO3- - 2SO, Reuss' and Johnson's definition of ANCaqua (easy to measure):
ANCaqua = 2Ca2+ + 2Mg2+ + K+ + Na+ - Cl- -NO3- -2SO^-
terrestrial plants, oceans have taken up considerable parts of it. A concern of many scientists is that rising levels of atmospheric CO2 are causing an increasing acidification of oceans. The latter is due to the equilibrium between atmospheric CO2 and the CO2 dissolved in seawater. Dissolved CO2 forms H2CO3 (Box 2), leading to increasing acidity. H+ reacts with the ion to form HCO^, that is, there is also a carbonate buffer in seawater.
CO2 is absorbed at the sea surface, which is thus most affected. Marine organisms that produce CaCO3 shells live above the so-called 'saturation horizon', where CaCO 3 does not readily dissolve. Increasing CO2 concentrations decrease the saturation state of CaCO3 making structures of CaCO3 vulnerable to dissolution.
Organisms containing shells or plates of CaCO3 fall to the sea floor when they die. CaCO3 is thus abundant in sediments and interacts with seawater. Slow mixing throughout the oceans, mixing necessary to bring up compounds from the oceans' sediments to buffer the increased ocean surface chemistry, causes a delay in CaCO3 dissolution. Warming of oceans as a consequence of global warming may even further reduce this mixing rate with deeper waters.
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