Ocean Acidification - "The Other CO2 Problem"
Blue Starfish resting on coral of the Great Barrier Reef off the coast of Queensland, Australia. Photograph by Richard Ling. Licensed under terms of the GNU Free Documentation License.
Ocean acidification, defined as the increasing hydrogen ion concentration in the world's oceans due to increases in atmospheric carbon dioxide (CO2) arising from human activities, can compromise the health and survival of many marine organisms. If continued at current rates, it threatens profound damage to the marine food web. The most vulnerable organisms are those that make exoskeletons containing calcium carbonate, but other organisms are also at risk.
When CO2 is added to the atmosphere, about one quarter of the added quantity dissolves in the ocean, and a roughly similar fraction is taken up by vegetation, soil, and other reservoirs on land, although the exact distribution between land and ocean is uncertain. In the ocean, the dissolved CO2 results in an increase in hydrogen ion concentration. This increase is conventionally expressed as a reduction in pH, which is the negative logarithm of the hydrogen ion concentration. The consequences for sea life are potentially far reaching, and ultimately pose threats to the welfare of millions of the world's population whose livelihood depends on the sea, and to the capacity of the oceans to serve as a source of food for human consumption.
Organisms at risk
The most vulnerable marine organisms are those that make calcium-containing exoskeletons (shells or shell like structures) important for their health and survival. The organisms include tiny plankton at the bottom of the food web, larger shellfish, a variety of algae, and corals whose reefs provide a haven for many species of fish. Of these, the corals may be the most conspicuous, but the plankton may be the most critical for the welfare of marine life overall. Although effects on calcification comprise the greatest threat, an increase in hydrogen ion concentration can also impair the internal metabolism of many marine organisms, including those that do not make calcified shells.
Ocean Carbonate Chemistry and Shell Calcification
The relevant chemistry can be described by a series of reactions that start with the combination of the dissolved CO2 with water to form carbonic acid - reaction : CO2 + H2O ↔ H2CO3. In turn, carbonic acid dissociates in water to yield a hydrogen ion and a bicarbonate ion - : H2CO3 ↔ H+ + HCO3-, and it is the resulting increase in ocean hydrogen ion concentration that defines the role of dissolved CO2 as a cause of ocean acidification. The bicarbonate can dissociate further to yield an additional hydrogen ion plus a carbonate ion - : HCO3- ↔ H+ + CO32-. In seawater, the tendency to form bicarbonate dominates, with this ion comprising almost 90% of the dissolved inorganic carbon derived from CO2, carbonate about 10%, and less than 1% remaining as CO2 and carbonic acid.
The link to calcified shells resides in the importance of carbonate concentration. To form solid calcium carbonate, the major shell component, calcium ions in seawater must be combined with carbonate ions. Although these ions can combine directly to form CaCO3, many marine organisms accomplish this indirectly, utilizing bicarbonate ions from seawater and from their own metabolism. The process involves a series of steps, yielding the following net reaction : Ca2+ + 2HCO3- → CaCO3 + H2O + CO2. It may seem paradoxical that the process of generating solid calcium carbonate also generates a CO2 molecule, but because two carbon atoms derivable from atmospheric CO2 are involved on the left side of the above reaction, and only one is released on the right side as CO2, the overall effect is the capture of a carbon atom in solid carbonate form as part of the structure of marine exoskeletal carbon. When the marine organisms die, some of their shells fall to the sea floor and are eventually compacted into limestone. This in turn is eventually removed from the climate system into the Earth's interior at subduction zones, and becomes part of a carbon cycle that returns the same carbon as CO2 emitted by volcanoes on land or the sea floor as well as by other sea floor vents for CO2.
Calcified Shells Can Dissolve
Organisms depending on calcified shells are subject to a major vulnerability - solid CaCO3 is susceptible to dissolution by reacting with water and CO2 to yield Ca2+ and a mixture of HCO3- and CO32- ions. Solid CaCO3 can only remain undissolved if Ca2+ and CO32- ion concentrations in seawater are high enough to maintain saturation; if the carbonate concentration declines sufficiently, the water becomes unsaturated, and solid CaCO3 continues to dissolve as long as unsaturation persists. Concentrations sufficient for saturation are defined by the solubility product constant, Ksp, whose magnitude is simply the value obtained by multiplying the Ca2+ and CO32- ion concentrations that are just high enough to achieve a saturated solution in which the ions are in equilibrium with solid CaCO3 - i.e, [Ca2+] x [CO32-]. Below those levels, solid CaCO3 will tend to dissolve. If ions are added to a level exceeding saturation, enough of them will combine to form solid CaCO3 to reduce their concentration to the saturation level. This may not occur spontaneously, but require the presence of a solid surface that can serve as a nucleus for crystal growth; in the absence of such a nucleus, the solution may remain supersaturated. (Technically, saturation is determined not by the ion concentrations per se, but by their "activity", representing an "effective concentration", but actual concentrations provide a good approximation). In seawater, Ca2+ is present in large excess, and saturation conditions are therefore largely a function of CO32- concentrations.
Ksp is an intrinsic property of different compounds, and even different forms of the same compound - for example, aragonite, the form of CaCO3 produced by corals, has a higher Ksp (i.e, is more soluble) than is a different type of crystalline CaCO3, calcite, which can persist in solid form at lower concentrations of Ca2+ and CO32-. Although referred to as a constant, Ksp varies with temperature, pressure, and salinity. Some of this variability is reflected in the observation that CaCO3 is more soluble at the high pressures and colder temperatures of the deep ocean than it is near the surface, and more soluble at colder latitudes than in the tropics. The ocean depth below which the water is sufficiently unsaturated for CaCO3 dissolution rates to increase dramatically is termed the lysocline.
An additional factor critical to the solubility of CaCO3 is the concentration of dissolved CO2. Because cold water can hold more CO2 in solution, the increase in CaCO3 solubility at colder temperature is explained at least in part by the higher CO2 concentration in the colder seawater. The effect of CO2 on CaCO3 solubility is explored further in the next section.
Effects of Rising CO2
When additional CO2 dissolves in seawater to yield carbonic acid (reaction  above), the predominant effect on the carbonate/bicarbonate balance is for hydrogen ions to combine with carbonate to form bicarbonate, thereby reducing the carbonate concentration - the net reaction  can be written as CO2 + H2O + CO32- ↔ 2HCO3-. If the carbonate concentration drops below saturation levels, the resulting dissolution of CaCO3 impairs the ability of calcifying organisms to manufacture their shells, and may be lethal unless the organisms can produce CaCO3 fast enough to outweigh the losses due to dissolution. Some organisms compensate by using their own metabolism to maintain adequate ion concentrations in their immediate vicinity despite unsaturated conditions in the surroundings, and a few even thrive at elevated CO2 and reduced pH, including photosynthesizing species that benefit from the additional CO2. On the other hand, many organisms, including many corals, require their surroundings to be supersaturated with Ca2+ and CO32- to enable them to produce sufficient shell material. This need arises because under natural conditions, their shells are subject to damage by erosion and predators, and so CaCO3 production must be rapid to keep pace with the losses. In addition, reduced pH environments may impose metabolic stress on calcifying organisms that impairs their ability to build shells, a process demanding high energy expenditure. The ion concentrations needed to balance all these threats are typically denoted by the Greek letter omega, Ω, where Ω = [Ca2+] x [CO32]/Ksp. A value of 1 signifies a saturating concentration, but under some stresses in the wild, a value of Ω closer to 3 may be needed for optimal calcification. At this level of supersaturation, a CaCO3 shell can efficiently serve as a nucleation site on which further CaCO3 can deposit to compensate for attrition. In laboratory experiments under protected conditions, the same organisms can often tolerate lower values for pH and Ω.
Photomicrographs of the shell of the pteropod, Clio pyramidata, collected from the subarctic Pacific. (a) Whole shell from a live pteropod kept in undersaturated seawater for 48 hours; the white rectangle indicates the location of the magnified area in (b), which shows advanced dissolution along the leading edge of the shell. (c) No dissolution is observed at the leading edge of shell from Clio pyramidata kept in seawater supersaturated with aragonite (photos from V. Fabry). Figure source: NOAA (National Oceanic and Atmospheric Administration)
Evidence from the Past
Preindustrial ocean pH averaged slightly more than 8.2, but over the course of more than a century of rising atmospheric CO2, it has declined to approximately 8.1, representing about a 30 percent increase in hydrogen ion concentration. Although this pH is still on the basic side of neutral, it is the increase in hydrogen ion that is biologically significant and which has earned the phenomenon the designation of ocean acidification. If CO2 continues to rise unabated during the remainder of the 21st century, reaching a level between 600 and 700 parts per million (ppm), an estimated further reduction to a pH level around 7.9 is considered probable; this would constitute an approximate doubling of hydrogen ion concentration over preindustrial levels, with profound effects on marine biology. A modest further rise in CO2 would most likely double hydrogen ion concentrations above current levels and reduce pH to about 7.8. .Estimates based on boron isotopes indicate that the current pH of about 8.1 is lower than at almost any time during the past 20 million years. A pH of 7.8 would be unmatched during the past 40 million years.
At this stage, it is difficult to discern how much harm, if any, has already resulted from the existing hydrogen ion increase. In the case of coral reefs, which are among the more conspicuous potential victims, it is likely that more harm has resulted to date from other influences, both natural (disease and predators) and of human origin (pollution and destructive fishing practices). Early warning signs have emerged, however, from studies on the Great Barrier Reef off the coast of Australia, where measurements have demonstrated a significant decline in calcification rates over the course of recent decades. This has not yet been matched by evidence of reef destruction, but evidence from the deep past tells us that a continued decline in pH is likely to exert a devastating influence on some corals that may require millions of years for recovery, and from which some areas may never recover. The extinction of other marine species during past intervals of low pH implies that many other organisms are also vulnerable.
Potential Impact of Rising Atmospheric CO2 on Coral Reef Calcification - figure source the National Oceanic and Atmospheric Administration (NOAA). Note that in 2011, CO2 levels had risen to 390 ppm.
Although calcifying organisms are the most vulnerable to increases in ocean acidity, changes in pH can also affect the internal metabolism of many other marine organisms as well, requiring them to spend metabolic energy to restore adequate function, and impairing their reproductive fitness. How much impact this phenomenon has already exerted is unknown. Additionally for corals, and probably other creatures as well, the warming effects of increased atmospheric CO2 can superimpose stresses on those created by the pH changes. "Bleaching" of coral reefs subjected to temperature stress (most often excessive heat but sometimes cold) has been one of the hallmarks of this phenomenon, and signifies that photosynthesizing organisms that exist in symbiosis with the corals have been lost or damaged. These symbionts are the source of most of the energy used by corals for calcification and other functions, and their loss, if prolonged, is often lethal.
Implications for the Future
Like its future effect on global temperature, the effect of CO2 on ocean pH will depend on the magnitude of future rises in atmospheric CO2 resulting from human activity. Although global warming has received more attention, it is not unlikely that the consequences of these separate CO2 effects will be comparable in severity. In particular, attempts to use geoengineering methods to ameliorate warming without reducing CO2 emissions would allow ocean acidification to proceed unabated, with potentially serious adverse effects on human welfare. In addition, measures designed to adapt to climate change are likely to be more feasible when applied to warming than to ocean acidification. Warming-induced rises in sea level can be addressed in some coastal areas by building sea walls and relocating at-risk populations, and agricultural threats from warming-induced drought can be partially ameliorated by expanded irrigation and the development of drought-resistant crops. It is harder, however, to conceive of practical measures that would avert much of the future harm that might ensue from widespread damage to the food web in the world's oceans. Reducing CO2 emissions remains the only clearly effective solution to both the warming and acidifying consequences of rising CO2.
Further Reading [with a brief summary of item content in brackets]
Fabry, V. et al, Impacts of ocean acidification on marine fauna and ecosystem processes, ICES Journal of Marine Science, 65: 414–432 (2008). [A comprehensive review and analysis of current and potential future effects of ocean acidification on marine organisms]
Pelejero, C. et al, Paleo-perspectives on ocean acidification. Trends in Ecology and Evolution, 25: 332-344 (2010). [Review briefly covering carbonate chemistry, effects on marine organisms, evidence from the distant past, and prospective future effects of further acidification]
Kroeker, K. et al, Meta-analysis reveals negative yet variable effects of ocean acidification on marine organisms. Ecology Letters 13:1419-1434 (2010) [Meta-analysis of 139 studies demonstrating variation in biological responses to ocean acidification, including survival, calcification, growth, and reproduction]
De'ath, G. et al, Declining Coral Calcification on the Great Barrier Reef. Science 323:116-119 (2009). [The title describes the content]
What Is Ocean Acidification? [An overview from NOAA]
Ocean Acidification [Overview and links from the U.S. National Academy of Sciences/National Research Council]
Frequently Asked Questions About Ocean Acidification [Answers to common questions about ocean acidification]