by Dr. Daniele Mazza
Oceans cover about 71% of the earth surface, but their influence on climate change is not only due to high heat capacity of water , not only to the ocean’s water circulation, but to a fact which is widely underestimated : the pH (acidity level) of sea-water is substantially alkaline, ranging from 8.0 to 8.7 . This means that the balance between positive and negative ions is reached by accounting for OH– ,hydroxide ions, in a far larger amount in respect to H+ hydrogen ions.
The pH value higher than 7 allows seawater to dissolve and react huge amounts of CO2 , carbon dioxide, thus affecting the amount of this gas in the atmosphere by absorbing excess of it. To calculate this excess in respect to what would be the true equilibrium value in the air, all of the chemical reactions involved have to be simultaneously computed, accounting for their equilibrium constants, which in turn depend on temperature.
1 – CO2 (gas) + H2O <==> H2CO3* (H2CO3* is the sum of dissolved CO2 and H2CO3)
2 – H2CO3 <==> H+ + HCO3–
3 – HCO3– <==> H+ + CO3– –
4 – H2O <==> H+ + OH–
5 – Ca++ + CO3– – <==> CaCO3 (calcite)
6 – Ca++ + OH– <==> Ca(OH)+
7 – Mg++ + OH– <==> Mg(OH)+
Before calculations, let us explore in some more detail mean seawater composition: summing up all the positive charges (Na+, K+, Mg++, Ca++) one obtains 621.1 moles per liter (mmol/L, or moles per cubic meter mol/m3). Carrying out the same operation for negative charges (Cl–, SO4– –, Br –) the result is slightly less : 619.2 mmol/L). 1.9 mmol/l are clearly missing ! The seawater must obey , as all other ionic solutions, to electrical neutrality law, so some negative ions have been ruled out: they are indeed HCO3– and to minor extent OH– and to far lesser extent CO3– –. All the last three ions are reactive, in respect of atmospheric CO2..
The presence of OH– ions (hydroxide ions) is the reason of a pH>7, their concentrations (due to the logarithmic nature of pH scale) is at pH = 8.0 equal to 0.001 mmol/L whilst that of H+ ions is 100 times less. OH– ions alone aren’t enough the fill the gap: we need other negative ions, these are mainly HCO3– ions, and also some CO3– – ions.
This fact has an immense consequence on the equilibrium of CO2 between atmosphere and oceans. Actual atmosphere contains around 850 Gt (giga tonn) of carbon (in form of CO2) while the oceans 38000 Gt of carbon, nearly 45 times more.
So when we talk about ppm CO2 in the atmosphere, that only is the top of the iceberg!
CO2 is a reactive gas, it dissolves (like N2 and O2) and later reacts with water itself (N2 and O2 do not) yielding HCO3– and CO3– – . After these reactions are completed still a third takes place (and is quite usually forgotten) : the formation of a solid salt, CaCO3 See reaction No 5 above. This is called in chemistry precipitation. CaCO3 usually has the form of calcite, aragonite, the other polymorph, is slightly more soluble. The seawater is oversaturated in respect of calcite, due to Ca++ ion concentration of 10.6 mmol/L . However this reaction require nucleation and growth of crystals and is usually sluggish (may speed up in the cell of invertebrates).
The destiny of this salt is to eventually sedimentate in the bottom of the sea, (may not reach the bottom, if very deep it can dissociate again in ions due to extreme high pressure and recycle again) . In any case the very end is to remove CO2 from the atmosphere forming limestone.
In textbooks of climate science or oceanography not always all the reaction are carefully accounted for the temperature influence.
Having taught applied chemistry at university level for more then 30 years, I found a simple but important point. When dealing with the above chemical equilibria, in most of the textbook, their equilibrium constant is considered constant, whilst these should vary with temperature.
I wrote some 300 line code in order to solve simultaneously all the above equilibria and to find if the actual level of 410 ppm of CO2 is in equilibrium or not with seawater carbonated ions. If not (and indeed it isn’t) how far are we from equilibrium and how does the system evolve in order to reach it?
Well I’ll try to resume, then if somebody is interested in detail, please e-mail me.
The complete list of considered equilibria is already written above, their equilibrium constants are calculated from Gibbs energy values (data are taken mainly from NIST database or other thermodynamic databases). Remember that K(eq) = exp(- ΔG/RT), R being the gas constants and T the absolute temperature.
1- ΔG = -20302 – T*(-96.25) (Joule/mol/K)
2- ΔG = 7660 – T*(-96.2) (Joule/mol/K)
3- ΔG = 14850 – T*(-148.1) (Joule/mol/K)
4- ΔG = 55836 – T*(-80.66) (Joule/mol/K)
5- ΔG = -13050 – T*(-202.9) (Joule/mol/K)
6- ΔG = -7576 (Joule/mol/K)
7- ΔG = -14656 (Joule/mol/K)
From the above treatment of inorganic carbon chemistry in seawater and the simultaneous resolution of temperature-dependent equilibria, interesting results are obtained. They are presented in graphic form, for sake of simplicity.
Figure 1 shows how CO2 if far more soluble in alkaline waters, like seawater. Compare the red line (ocean water) with blue one (pure water). On x-axis are ppm CO2 in standard air at 17°C (from 200 to 600) and on y-axis the C(T) , total (inorganic) carbon content, i.e. the sum of CO2(aq) , H2CO3 , HCO3– and CO3– – . Note how, increasing ppm CO2 pH changes slightly from 8,72 to 8,27 not so dramatically.
Next figure 2 indicates how temperature affects the inorganic carbon equilibria at constant CO2 (400 ppm). With increasing temperature the DIC or C(T) (total dissolved inorganic carbon) decreases and pH increases.
This explains why CO2 is released in air in warm equatorial waters and absorbed in cold waters.
Figure 3 gives us a comprehensive view of the degree of non-equilibrium in the (average) CO2 exchange between air and ocean. The blue point represents actual 400 ppm value that should reach the 315 ppm equilibrium value with an average sea temperature of 17°C
Ocean water are therefore a huge reservoir for CO2 that waits to be filled.
Up to now calcite precipitation isn’t taken into account. But this is done in figure 4, which explains how CaCO3 forms and thus collects still other huge quantities of CO2. Red curve represents C(T) as a function of temperature with no precipitation of calcite (the same as fig. 1), green line with complete calcite precipitation (We assume C(T) = 1.85 to be a constant value).
The real situation is slowly moving from the red to the green curve, which will be reached at the end. How long does it take? Should be a question of some years or more but the phenomenon will go that way and not the reverse. Blue line is the quantity of limestone at the end of process (green curve). Limestone in geological time will be pushed to subduction by plate tectonic movements, heated by magma and in the far end decomposed to CO2 and calcium silicates. CO2 will be emitted in the volcanoes again in the air after million of years, far enough so that all fossil fuels are burned out !
Conclusions : CO2 is at 410 ppm far above the equilibrium value (315) , provided a standard seawater composition and an average ocean temperature of 17°C (taken from wikipedia). No doubt that solubility will force more CO2 to be stored in oceans . Moreover if we consider CaCO3 formation (seawater has overshot the solubility of this salt nearly 50 times but nucleation and growth are slow) still more CO2 will be stored by limestone.