Guest essay by James R. Barrante, Ph.D.
Emeritus Professor of Physical Chemistry, Southern Connecticut State University, New Haven, CT
It is well-known that ocean pH has decreased from approximately 8.2 pH units to 8.1 pH units in the last 150 years. One explanation is that the increase in atmospheric carbon dioxide from the burning of fossil fuels is responsible. This is based on a law of physical chemistry, known as Henry’s Law, that states: if the partial pressure of a gas over a solution is increased, the concentration of the dissolved gas in the solution will increase. Since dissolved CO2 is a weak acid, one would expect the pH of the oceans to decrease.
Describing Henry’s Law in this simple manner, however, is a very narrow interpretation of the boundaries of the system. For example, the above statement of Henry’s Law is only valid, if the temperature of the solution is constant. Moreover, using pH to describe the acidity or alkalinity of a solution can be misleading. The pH of a pure water solution (neither acidic nor alkaline) is 7.00 at 298.2 K. Increase the temperature of the water and the pH will drop to below 7.00. Does this mean that simply raising the temperature of water will cause water to become “acidic?” That would be a ridiculous interpretation of pH.
Thermodynamic equilibrium constants are a sensitive function of temperature. We know that if the temperature of an aqueous solution of a gas increases, the solubility of that gas has to decrease, not increase. Moreover, if the ionic strength of the solution is high, describing equilibrium constant equations in terms of concentration will produce, at best, only approximate results. Activities must be used. Ocean chemistry is complex, involving a number of important equilibria, that include the dissociation of carbonic acid (dissolved CO2), the buffering equilibria due to the presents of dissolved salts of bicarbonate and carbonate, the solubility of the sparingly soluble salt CaCO3, and the equilibrium between dissolved CO2 and the partial pressure of carbon dioxide in the atmosphere. Luckily, the temperature dependence of these equilibrium constants has been highly studied. A typical example of each is given below.
The pH of seawater (ionic strength of approximately 0.7) can be determined at any temperature a modified form of the Henderson-Hasselbalch equation for buffer systems (derived in my blog: climaterx.wordpress.com., post OAII). For reasons described in the post, the activity of calcium ion in the oceans is assumed to be constant.
To test this equation, consider the following:
Preindustrial pH: PCO2 = 0.000280 atm
aCa++ = 0.00123 (described in blog)
T = 288.2 K
pK1 = 6.4149; pK2 = 10.4202
kH = 22.24 atm/C
Ksp = 6.05 x 10-9 (average calcite/aragonite)
pH = 8.214 (very close to published value)
Post-industrial pH: PCO2 = 0.000380 atm
aCa++ = 0.00123 (described in blog)
T = 290.2 K
pK1 = 6.4105; pK2 = 10.4063
kH = 23.56 atm/C
Ksp = 5.718 x 10-9 (average calcite/aragonite)
pH = 8.138 (very close to published value)
We find from similar calculations that a 2-degree C increase in temperature can lower ocean pH by approximately 0.05 pH units. The temperature range taken in this study was ocean SST at the Equator, approximately 305 K, to ocean SST at the poles, approximately 273 K. This alone with play havoc with pH measurements. Moreover, ocean temperature ranges in bands containing water of different surface areas and land masses running parallel to the Equator. Simply recording ocean temperature at various points over the ocean and averaging them to get a number has no useful scientific meaning, no more than the average diameter of a football tells us anything about the shape of a football. At best, a weighted average must be used. In fact, if the measurement of ocean pH were not so complicated, and we had that data for the last 150 years, I would bet that we could show exactly that the increase in atmospheric CO2 from 280 ppmv to 380 ppmv in the last 150 years is an ocean temperature effect and not at all related to burning fossil fuels.
The graph below is a three-dimensional representation of P-pH-T data, similar to PVT graphs. The two-dimensional dependence of pH on either CO2 pressure or ocean surface temperature can be followed by looking at various isotherms or isobars; however, be sure to follow the graphs back to their original axes. For convenience, pH values at the intersection of these curves are given. It is obvious that to assume the drop in alkalinity of such a highly buffered system as our oceans is due to an uptake of atmospheric CO2 without taking the increase in ocean temperature, particularly in the Northern Hemisphere, into account is a very narrow interpretation of the science.
Temperature and Pressure Behavior of Ocean pH
Note: the word “only” was added to the title for clarity – Anthony