ACS Climate Science Toolkit | Oceans, Ice, and Rocks
Take a deep breath. It’s possible that among the molecules of the CO2 you inhaled some were emitted from burning the first tonne of coal in the first steam engine built by James Watt—the engines that powered the Industrial Revolution. Since that first engine began operating, more than 300 gigatonnes of carbon, GtC (≈ 1,100 Gt CO2), from the fossil fuel reservoir in this figure have been added to the atmosphere by human activities. About two-thirds is still there, included in the 800 GtC in the figure. Much of the rest is in the ocean.
Assuming that about half of Watt’s emitted CO2 is present and well-mixed in the atmosphere, one breath would contain over 7 million of those molecules. This neglects the exchanges with other reservoirs shown in the figure.
Chemistry, Atmospheric CO2, and Ocean Acidification,” Annu. Rev. Earth Planet. Sci.,
2012, 40, 141-165
In addition to the ocean’s role as an energy reservoir moderating Earth’s climate, this figure illustrates its important role as a carbon reservoir interacting with the atmosphere. Almost all of the carbon in the atmosphere, the ocean, and the sediments is in its most oxidized form as CO2 or carbonates, HCO3– and CO32–, in solution or precipitates, largely CaCO3. The interactions shown in the figure had been essentially in equilibrium for millennia before the Industrial Revolution. The sudden (on a geological timescale), and continuing, two-century injection of CO2 into this chemical system has upset the balance and we are now observing the effects of the disturbed annual to centennial interchanges noted in the figure. To see how adding CO2 affects the system, first we need to consider the major reactions going on in the ocean-atmosphere system.
This figure represents how carbon is stored and interchanged among various reservoirs on Earth. The numbers in the labeled boxes are gigatonnes (1015 g) of carbon, GtC, in that reservoir. The numbers associated with the interchange arrows are approximate annual fluxes, GtC·yr–1, between reservoirs. The numeric values are for 2012. The timescales are an indication of how long it takes a change within the denoted reservoirs to come to equilibrium. Sedimentary rocks, formed from these sediments deep in the Earth over even longer times than here, constitute the largest carbon reservoir of all, about 100,000,000 GtC.
The solubility of CO2 in seawater increases with the pressure of the gas and decreases as the temperature and salinity of the seawater increase. The dissolution equilibrium reaction is
CO2(g) ⇔ CO2(aq). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . (1)
Dissolution of gases in the ocean takes a relatively long time to come to equilibrium. Mixing gas and liquid at the surface is inefficient, even with wind agitating the liquid surface. Mixing of the upper layers of the ocean with the next lower layers takes many years, as exemplified by energy transfer in the top layers, and movement of surface layers to the deep ocean takes centuries. Thus, some of the CO2 from fossil fuel combustion dissolves in the top layers of the ocean in a few years, while some of the rest requires centuries or millennia to come to solubility equilibrium with the entire ocean.
This equilibrium is temperature sensitive. For example, during a glacial period, the system has many thousands of years to reach equilibrium at a low ocean temperature. As the ocean temperature increases at the beginning of an interglacial period, the equilibrium shifts to the left. CO2 becomes less soluble, its atmospheric concentration increases, and it contributes to further warming via its greenhouse effect.
Once dissolved, the rapid reactions of CO2(aq) with water and further acid-base and ionic solubility reactions are what make the system more complex than the simple dissolution of a nonreactive gas.
CO2(aq) + H2O ⇔ HCO3–(aq) + H+(aq) . . . . . . . . . . . . Ka1 . . . . . . . . . . . . . . . . (2)
HCO3–(aq) ⇔ CO32–(aq) + H+(aq) . . . . . . . . . . . . . . . . . . Ka2 . . . . . . . . . . . . . . . . (3)
The equilibrium constants here represent the successive acid dissociation constants for carbonic acid, H2CO3. However, aqueous solutions of CO2 contain only a minuscule amount of H2CO3, so the production of the initial hydronium ion, represented by H+(aq), is more accurately characterized by equation (2) and its equilibrium constant.
The Arrhenius acid-base formulation is used here, even though solvated protons, H+(aq), do not exist in aqueous solution, because it helps to focus on the carbon-containing species. Explicitly including the hydronium ion, H3O+, and the Brönsted-Lowry base, H2O, needlessly complicates the equations.
In seawater,at pH = 8.2, the distribution among the three inorganic carbon-containing species, CO2(aq), HCO3–(aq), and CO32–(aq), is about 0.5%, 89%, and 10.5%, respectively. You see that the first equilibrium, equation (2), lies far to the product side, with very little CO2(aq) left unreacted in solution, while the second, equation (3), favors the reactant side, thus making the bicarbonate ion, HCO3–(aq), the predominant inorganic carbon species in the ocean.
The ultimate fate of this inorganic carbon (that does not get incorporated into organic carbon via photosynthesis) is the formation of insoluble ionic salts, the great majority of which is CaCO3. To help understand the effect of adding more CO2 to the ocean, we write the formation of CaCO3 this way
Inorganic carbon in the ocean is the carbon in the three species here and the ionic compounds formed from them. Organic carbon is the carbon in molecules that are synthesized by living things, beginning with photosynthesis using carbon from the inorganic precursor, CO2. This includes molecules in the living things—phytoplankton, krill, and whales, etc.—or in their remains. Like inorganic carbon, most organic carbon ultimately ends up in sediments on the ocean floor.
Ca2+(aq) + 2 HCO3–(aq) ⇔ CaCO3 + CO2(aq) + H2O. . . . . . . . . . . . . . . . . . . . . . (4)
Although CaCO3 is insoluble, the formation of the precipitate at the concentrations and salinity in seawater is slow. The solutions tend to supersaturate, so, in this equation, the ionic salt is not characterized as either solid or aquated.
Many marine organisms use CaCO3 as part of their structure. Corals build some of the most striking of these structures, including coral reefs, like the one pictured here. Corals are beautiful ecotourism attractions and provide homes to the greatest biodiversity in the ocean, but they occupy only a tiny part of the ocean in shallow waters near land and have little effect on Earth’s carbon budget.
The reactions in equation (4) are easy to demonstrate. Mixing concentrated solutions of CaCl2 and NaHCO3 produces immediate copious bubbling and froth together with the formation of a white precipitate. Bubbling CO2 gas into limewater, aqueous Ca(OH)2, first produces a white precipitate of CaCO3(s), which dissolves as CO2 continues to be bubbled into the suspension.
Microscopic organisms, represented by the coccolithophore phytoplankton in the scanning electron micrograph below are found in the surface of all the oceans across the globe and make a major contribution to Earth’s carbon budget. These are photosynthetic organisms, so they remove CO2 from the ocean-atmosphere system as they produce the organic molecules that sustain their life. They also take in Ca2+(aq) and HCO3–(aq) ions from the surrounding sea to synthesize the intricate CaCO3 scales that adorn their exterior. In the process of making CaCO3, they produce CO2, as equation (4) shows. As living organisms, it is not certain whether the coccolithophores represent a net sink or source for CO2.
There is no doubt, however, that, taken together, coccolithophores represent the largest source of biogenic CaCO3 on Earth, contributing significantly to the global carbon budget. The scales they synthesize and shed throughout their lives and those remaining when they die either redissolve or add to the sediment at the bottom of the oceans. Their fossil remains are found in these sediments as well as in the great limestone deposits, such as the White Cliffs of Dover, that form from these sediments after millions of years.
Effect of Added Atmospheric CO2: Ocean Acidification
As the amount of CO2 in the atmosphere increases, Le Chatelier’s principle applied to equation (1) shows that, CO2(aq), the amount in the oceans, also increases. Le Chatelier’s principle applied to reaction (2) shows that increasing CO2(aq) favors the forward reaction, increasing the HCO3–(aq) and H+(aq) concentrations. Similarly, Le Chatelier’s principle applied to reaction (4) shows that increasing CO2(aq) favors the reverse reaction, increasing Ca2+(aq) and HCO3–(aq), while decreasing CaCO3 concentrations. CO2(aq) is acidic, so reacts with the strongest base present in significant concentration, CO32–(aq), and upsets the carbonate buffer system.
CO2(aq) + CO32–(aq) + H2O ⇔ 2 HCO3–(aq) . . . . . . . . . . . . . . . . . . . . . . . . . . (5)
The net result of these shifts in equilibrium concentrations is increased H+(aq), a more acidic ocean, and decreased availability of CaCO3 and CO32–(aq). The experimental observation is that the pH of the top layer of the ocean has decreased by 0.1 unit, from 8.2 to 8.1 during the past century or so as about 100 GtC (≈ 370 Gt CO2) have dissolved in the oceans. As more CO2 dissolves, these concentration changes will continue and effects on marine life become evident. Examples might include the displacement of some species by others better adapted to a more acidic environment or the inability of some shell-forming organisms to survive in an environment that favors dissolution of CaCO3. The latter is of great concern because many of these organisms are at the base of the oceanic food chain that furnishes a substantial percentage of the human food supply. The extent and direction of the effects of a changing ocean are not yet certain, so a great deal of effort is being expended to find out what they are and will be.