Acidification of Water by CO2
Acidification of Water by CO2
W. A. van Wijngaarden
Department of Physics and Astronomy, York University, Canada
P. Ridd
Physicist, Adjunct Fellow, Institute of Public Affairs, Malanda, Queensland, Australia
M. Cornell
Industrial Chemist, Lake Jackson, Texas, USA
W. Happer
Abstract
Fundamental inorganic chemistry shows that increasing concentrations of atmospheric CO2 will have no harmful effect on organisms that live in the natural waters of the Earth [1], and may well benefit them. Alkalinity and dissolved CO2 give high buffering capacity to most natural waters and minimize the change of pH from external influences. For example, doubling the atmospheric concentration of CO2 from 430 ppm to 860 ppm would reduce the pH of representative sea water at a temperature of 25 C from pH = 8.18 to pH = 7.93. This change is comparable to diurnal pH changes in biologically productive surface waters, due to photosynthetic fixation of dissolved inorganic carbon during the day and respiration at night. The change is also less than the variations of pH with latitude, longitude and depth in the oceans. This paper includes a quantitative review of the carbonate chemistry of seawater and freshwater, the buffering capacity, the Revelle factor, the transport of calcium carbonate in ground water, the formation of flowstone, and the classic use of limewater to detect gaseous CO2. The paper concludes with a brief review of those parts of chemical thermodynamics that are involved in ocean acidification.
Keywords: ocean acidification, pH, alkalinity, buffering, carbonate chemistry, groundwater
1 Introduction
Natural surface waters and rainwater contain dissolved CO2, in equilibrium with atmospheric CO2. As atmospheric CO2 concentrations increase, the increase of the dissolved weak acid, CO2, will lower the pH of the waters. Fig. 1 shows how the pH of representative seawater and rainwater depend on the partial pressure P of CO2 [1].
The lowering of the pH of natural surface waters by more atmospheric CO2 is often called acidification, but this term can be misleading. For example, representative seawater at a temperature T=25 C is quite basic, with a pH = 8.18 at current partial pressures, Pc=430 μb of CO2. Doubling the partial pressure of CO2 to Pd=860 μb would decrease the pH of the ocean to pH = 7.93, a decrease of ΔpH=−0.25. This would leave the oceans almost as basic as today. At the same temperature, T=25 C, doubling of CO2 would reduce the pH of rainwater, which is slightly acidic because of dissolved CO2 and with no compensating alkalinity, from pH = 5.59 to pH = 5.44, an even smaller decrease of ΔpH=−0.15. Although contemporary seawater has absorbed about 120 times more CO2 per unit volume than rainwater in equilibrium with the same partial pressure P of CO2, the relatively large alkalinity [A] = 2.4 mM of seawater keeps it basic.
In subsequent sections we will show straightforward ways to calculate the numbers cited here. Lord Kelvin once said [2]:
“When you can measure what you are speaking about, and express it in numbers, you know something about it, when you cannot express it in numbers, your knowledge is of a meager and unsatisfactory kind; it may be the beginning of knowledge, but you have scarcely, in your thoughts advanced to the stage of science.”
As Lord Kelvin urged, much of the discussion below is in terms of numbers, equations and graphs. We have simplified discussions of the complicated chemistry of the oceans and the biochemistry of ocean life to keep the paper understandable to readers who recall some high school mathematics, chemistry and biology. More details can be found in the literature cited.
The oceans will remain non-acidic for any credible scenarios of human emissions of CO2 in the future. This paper gives a quantitative discussion of why increasing concentrations of atmospheric CO2 will not cause harmful acidification of natural waters on Earth, and will have little effect on the biochemistry of aquatic organisms.
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