4. Chemical Laws for Distribution of CO₂ in Nature

Volk & Liu (1988) modelled the CO₂ flux between atmosphere and oceans, and concluded that approximately 70% of the flux was governed by this "thermal solubility pump", while approximately 30% was governed by the organic nutrient "biological pump". Faure (1990) estimated that ca. 4000 GT (Gigatonnes = billion metric tonnes) of CO₂ is transferred by degassing of the ocean via the atmosphere to the continental biosphere from the end of a glaciation to an interglacial stage.

From a geochemical consideration of sedimentary rocks deposited throughout the Earth's history, and the chemical composition of the ocean and atmosphere, Holland (1984) showed that degassing from the Earth's interior has given us chloride in the ocean; and nitrogen, CO₂, and noble gases in the atmosphere. Mineral equilibria have established concentrations of major cations and H⁺ in the ocean, and the CO₂ concentration in the atmosphere, through different chemical buffer reactions. Biological reactions have given us sulphate in the ocean and oxygen in the atmosphere.

Carbon dioxide is an equally important requisite for life on Earth as oxygen. Plants need CO₂ for their living (the photo synthesis), and humans and animals breath out CO₂ from their respiration. In addition to this biogeochemical balance, there is also an important geochemical balance. CO₂ in the atmosphere is in equilibrium with carbonic acid dissolved in the ocean, which in term is close to CaCO₃ saturation and in equilibrium with carbonate shells of organisms and lime (calcium carbonate; limestone) in the ocean through the following reactions (where s indicates the solid state, aq is the aqueous state, and g is the gaseous state):

Partial reactions:

CO₂ (g)    <=>    CO₂ (aq)

CO₂ (aq) + H₂O    <=>    H₂CO₃ (aq)

H₂CO₃ (aq)    <=>    H⁺ (aq) + HCO₃^- (aq)

HCO₃^- (aq)    <=>    H⁺ (aq) + CO₃^2- (aq)

CO₃^2- (aq) + Ca²⁺ (aq)    <=>    CaCO₃ (s)

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Net reaction:

CO₂ (g) + H₂O + Ca²⁺ (aq)    <=>    CaCO₃ (s) + 2 H⁺ (aq)

In addition there are a number of different aqueous metal complexes of lesser concentrations.

A buffer can be defined as a reaction system which modifies or controls the value of an intensive (i.e. mass independent) thermodynamic variable (pressure, temperature, concentration, pH, etc.). Our carbonate system above will act as a pH buffer, by the presence of a weak acid (H₂CO₃) and a salt of the acid (CaCO₃). The concentration of CO₂ (g) and of Ca²⁺ (aq) will in the equilibrium Earth system also be buffered by the presence of CaCO₃, at a given temperature. If the partial pressure of CO₂ (g) is increased, the net reaction will go towards the right because of the Law of Mass Action. If the temperature changes, the chemical equilibrium constant will change, and move the equilibrium to the left or right. The result is that the partial pressure of CO₂ (g) will increase or decrease. The equilibrium will mainly be governed by Henry's Law: the partial pressure of CO₂ (g) in the air will be proportional to the concentration of CO₂ (aq) dissolved in water. The proportional constant is the Henry's Law Constant, which is strongly temperature dependent, and lesser dependent on total pressure and salinity (Drummond, 1981).

Questions have been raised about how strong this buffer is. It has been postulated (Bolin & Keeling, 1963) that an increase in atmospheric CO₂ will be balanced when only approximately one tenth of this is dissolved in the ocean. This postulate fails for a number of reasons. An increase in atmospheric CO₂ will namely increase the buffer capacity of ocean water, and thereby strengthen the ocean's capacity to moderate an increase of atmospheric CO₂; maximum buffer capacity for the system CO₂ - H₂O is reached at 2.5 to 6 times the present atmospheric partial pressure of CO₂, depending on temperature and alkalinity (Butler, 1982). According to Maier-Reimer & Hasselmann (1987) the borate system also increases the ocean storage capacity for CO₂ by more than 20% over an ocean with the carbonate-system alone.

Furthermore, this carbonate buffer is not the only buffer active in the atmosphere / hydrosphere / lithosphere system. The Earth has a set of other buffering mineral reactions. The geochemical equilibrium system anorthite CaAl₂Si₂O₈ - kaolinite Al₂Si₂O₅(OH)₄ has by the pH of ocean water a buffer capacity which is thousand times larger than a 0.001 M carbonate solution (Stumm & Morgan, 1970). In addition we have clay mineral buffers, and a calcium silicate + CO₂ <-> calcium carbonate + SiO₂ buffer (MacIntyre, 1970; Krauskopf, 1979). These buffers all act as a "security net" under the most important buffer: CO₂ (g) <-> HCO₃^- (aq) <-> CaCO₃ (s). All together these buffers give in principle an infinite buffer capacity (Stumm & Morgan, 1970).

Stable carbon isotopes (¹³C/¹²C) show that CO₂ in the atmosphere is in chemical equilibrium with ocean bicarbonate and lithospheric carbonate (Ohmoto, 1986). The chemical equilibrium constants for the chemical reactions above provide us with a partition coefficient for CO₂ between the atmosphere and the ocean of approximately 1 : 50 (approx. 0.02) at the global mean temperature (Revelle & Suess, 1957; Skirrow, 1975). This means that for an atmospheric doubling of CO₂, there will have to be supplied 50 times more CO₂ to the ocean to obtain chemical equilibrium. This total of 51 times the present amount of atmospheric CO₂ carbon is more than the known reserves of fossil carbon. It is possible to exploit approximately 7000 GT of fossil carbon, which means, if all this carbon is supposed to be burned, that the atmospheric CO₂can be increased by 20% at the most under geochemical equilibrium at constant present surface temperature.

¹⁴C isotopes show that the circulation time for carbon in the upper part of the ocean is some few decades (Druffel & Williams, 1990). This is sufficient time for the ocean to absorb an increase in atmospheric CO₂ from burning of fossil fuel at the present projected rate (Jaworowski et al., 1992 a).

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