Research Paper On Ocean Acidification Coral Reefs

Abstract

Ocean acidification is likely to have direct negative physiological consequences for many marine organisms, and cause indirect effects on marine ecosystems. Ocean acidification could also affect the oceans' current role as a net carbon sink by altering the oceanic calcium carbonate budget. Although ocean acidification and climate change are both caused by greenhouse gas emissions, ocean acidification is not climate change per se, and is often referred to as “the other carbon dioxide (CO2) problem.” As the United States considers actions in response to climate change, it is critical to take into account not only the impact of CO2 emissions on the climate but also their ramifications for ocean chemistry. The metrics that currently guide the climate change debate are dominated by strategies to reduce thermal impacts on the terrestrial environment. In this article, I examine the effects of ocean acidification and why they should help guide decisionmakers in setting CO2 emissions goals.

The oceans have absorbed approximately 30% of the atmospheric carbon dioxide (CO2) emitted by humankind since the beginning of the Industrial Revolution, and they serve as an important net carbon sink (Sabine et al. 2004). Although oceans, as carbon sinks, reduce the rate of CO2 increase in the atmosphere, the absorption process has a direct and measurable impact on ocean chemistry. Since the late 18th century, the mean surface ocean pH has dropped 0.1 units from 8.2 to 8.1. Because pH is measured on a logarithmic scale (a unit change of 1.0 is equal to a tenfold change in concentration), this change is roughly equivalent to a 30% increase in the concentration of hydrogen ions (Raven et al. 2005). Unlike changes in temperature due to global warming, which are difficult to predict, the mean magnitude and rate of ocean acidification can be projected with high confidence under different CO2 emissions scenarios, derived from a series of highly predictable chemical reactions (Caldeira and Wickett 2003, Raven et al. 2005). Under the Intergovernmental Panel on Climate Change (IPCC) business-as-usual (BAU) IS92a scenario, in which atmospheric concentrations of CO2 are expected to reach 800 parts per million by volume (PPMV) by 2100 (current atmospheric carbon is around 384 PPMV), the mean surface ocean pH is projected to drop another 0.3 to 0.4 units (Orr et al. 2005)—a 150% increase in the concentration of hydrogen ions.

When CO2 dissolves in seawater, it generates carbonic acid (H2CO3), which breaks down into bicarbonate (HCO3), carbonate (CO32−), and hydrogen ions (H+; figure 1). These reactions are reversible and near equilibrium (Millero et al. 2002). Their directionality is determined by ion concentrations (including noncarbon ions), temperature, salinity, and pressure. Under the oceans' current physical and biological conditions, carbonate and bicarbonate ions act as a buffer by absorbing and storing excess carbonic acid. This process buffers seawater against larger changes in pH. Under high-CO2 conditions, the reactions forming hydrogen ions and bicarbonate are favored, and the result is lowered pH and fewer carbonate ions (figure 2).

Figure 1.

Seawater carbonate chemistry equations. Carbon dioxide dissolves into seawater from the atmosphere and generates carbonic acid (H2CO3), which breaks down into bicarbonate (HCO3), carbonate (CO32−), and hydrogen ions (H+). When protons combine with carbonate ions to form bicarbonate, the concentration of carbonate decreases, making it unavailable to marine calcifiers to form calcium carbonate (CaCO3). All of these reactions are reversible, and directionality depends on concentration, temperature, salinity, and pressure. Source: Adapted from Hoegh-Guldberg et al. 2007.

Figure 1.

Seawater carbonate chemistry equations. Carbon dioxide dissolves into seawater from the atmosphere and generates carbonic acid (H2CO3), which breaks down into bicarbonate (HCO3), carbonate (CO32−), and hydrogen ions (H+). When protons combine with carbonate ions to form bicarbonate, the concentration of carbonate decreases, making it unavailable to marine calcifiers to form calcium carbonate (CaCO3). All of these reactions are reversible, and directionality depends on concentration, temperature, salinity, and pressure. Source: Adapted from Hoegh-Guldberg et al. 2007.

Figure 2.

Bjerrum plot. The relative proportions of carbonate species concentrations (in moles per kilogram) vary as a function of pH. Depicted here is a hypothetical shift in mean surface water pH from 8.2 (preindustrial) to 7.75, the value at which calcium carbonate becomes undersaturated with respect to aragonite. Under high-CO2 conditions, the reactions forming hydrogen (H+) and bicarbonate (HCO3) are favored, and the result is fewer carbonate ions (CO32+).

Figure 2.

Bjerrum plot. The relative proportions of carbonate species concentrations (in moles per kilogram) vary as a function of pH. Depicted here is a hypothetical shift in mean surface water pH from 8.2 (preindustrial) to 7.75, the value at which calcium carbonate becomes undersaturated with respect to aragonite. Under high-CO2 conditions, the reactions forming hydrogen (H+) and bicarbonate (HCO3) are favored, and the result is fewer carbonate ions (CO32+).

A paleoclimatic event relevant to the current ocean situation occurred 55 million years ago when the atmosphere had higher concentrations of greenhouse gases and a higher mean global temperature. It has been hypothesized that there was an input of CO2 to the deep ocean, presumably the result of a massive volcanic methane release of about 2000 gigatons of carbon over approximately 10,000 years (Zachos et al. 2005, 2008). In this scenario, methane would have been rapidly oxidized into CO2, lowering the pH and carbonate ion concentration in the deep sea. (For reference, since the beginning of the Industrial Revolution, humans will have released 5000 gigatons by 2400 under the IS92a IPCC BAU scenario [Caldeira and Wickett 2003]). The fossil record indicates that this massive methane-release event was followed by the extinction of several bottom-dwelling foraminifera species, single-celled organisms with calcareous shells (Zachos et al. 2005). Although knowledge of this ancient event may help us to predict what might happen to ocean pH and biota under contemporary conditions (Zachos et al. 2005), today's world is much different. Current rates of ocean acidification are faster, preindustrial levels of CO2 and temperature were lower, and a vastly different marine biota occupies the oceans today (Doney et al. 2009).

Impacts of ocean acidification on marine biota

Changes in ocean chemistry will probably affect marine life in three different ways: (1) decreased carbonate ion concentration could affect the calcification process for calcifying organisms (e.g., corals); (2) lowered pH could affect acid-base regulation, as well as a variety of other physiological processes; and (3) increased dissolved CO2 could alter the ability of primary producers to photosynthesize. Most of the research in the field has focused on calcification effects. In this article, I focus on how the concentration of carbonate ions affects calcification and dissolution, and then summarize what is known about other physiological effects, including those relating to photosynthesis. Following these species-specific responses, I provide an overview of theoretical community and ecosystem effects that might be expected into the future.

The importance of the carbonate ion for calcification. Many organisms use calcium and carbonate ions from seawater to produce calcium carbonate (figure 1), a compound used for skeletal support (e.g., corals) and protection (e.g., snail shells). The three mineral forms of calcium carbonate commonly produced are aragonite, calcite, and high-magnesium calcite (Raven et al. 2005). Aragonite calcifiers (e.g., stony corals and shelled pteropods) and high-magnesium calcite calcifiers (e.g., coralline algae and sea urchins) are likely to be affected by ocean acidification more strongly than calcite calcifiers (e.g., foraminifera and coccolithophores). This distinction is because of differences in the solubility of mineral forms; for example, aragonite is approximately 50% more soluble than calcite (Doney et al. 2009).

The saturation state, denoted by the Greek letter Ω, refers to the degree to which seawater is saturated with a carbonate mineral (i.e., aragonite, calcite, or high-magnesium calcite), and is inversely related to the mineral's solubility. The saturation state is determined by the concentrations of calcium and carbonate ions in relation to the solubility coefficient (K'sp) for the particular calcium carbonate mineral (Ω = [Ca2+][CO32−]/K'sp). The solubility coefficient varies with temperature, salinity, and pressure. Calcium carbonate solubility rises with decreasing temperature and increasing pressure and therefore increases with ocean depth. Since the saturation state and the solubility coefficient are inversely related for a given ion concentration, the saturation state is highest in shallow, warm tropical waters and lowest in deep and cold high-latitude waters (Feely et al. 2004).

When the saturation state is equal to 1 (Ω = 1), there is an equal chance of dissolution or formation of calcium carbonate. This chemical threshold defines a two-dimensional surface in the ocean interior called the saturation horizon. When saturation is greater than 1 (Ω > 1), formation of calcium carbonate is favored; when saturation is less than 1 (Ω < 1), dissolution is favored. The saturation horizon normally occurs at an ocean depth separating shallow, warm surface waters from deeper, cooler water. However, ocean acidification is decreasing the concentration of carbonate ions and therefore decreasing the saturation state of calcium carbonate, thus bringing the saturation horizon closer to the surface (Raven et al. 2005). Because of the higher solubility of aragonite and high-magnesium calcite, the saturation horizons for these mineral forms are even shallower than for calcite. Saturation horizons also vary in depth between ocean basins; aragonite and calcite saturation horizons are shallower in the Indian and Pacific oceans than in the Atlantic Ocean because of their longer deep-water circulation pathways, resulting in a greater accumulation of biologically respired CO2 (Broecker 2003).

For many organisms, a decrease in the calcium carbonate saturation state has been correlated with a reduction in calcification rates, even when waters remain supersaturated (Ω > 1) with respect to aragonite and calcite (Kleypas et al. 2005, Fabry et al. 2008). Under laboratory conditions, several calcifying organisms, including abundant planktonic species (e.g., coccolithophores, pteropods, foraminifera) and benthic invertebrates (e.g., coral, calcifying algae, molluscs, echinoderms), have shown a reduction in calcification rates as a result of reduced carbonate ion concentration (Raven et al. 2005, Fabry et al. 2008, Guinotte and Fabry 2008). However, the impact of lower carbonate concentration on calcification rates varies among species, and there is evidence that acidification may even enhance calcification in some taxa (Iglesias-Rodriguez et al. 2008, Ries et al. 2008, Wood et al. 2008).

A decrease in calcium carbonate saturation can also increase the dissolution rate of existing exposed and unprotected (i.e., by organic coatings) calcium carbonate structures (Jokiel et al. 2008, Andersson et al. 2009). As saturation horizons become shallower, more organisms will inhabit undersaturated waters (Ω < 1) where unprotected calcium carbonate structures will dissolve. The rate and extent of this process varies across the oceans because the solubility coefficient depends in part on temperature, pressure, and salinity (Raven et al. 2005).

Other physiological effects. Other potential negative physiological effects of ocean acidification include impairment of acid-base regulation, reproduction (sea urchins), respiration (squid, crabs), metabolism (mussels, worms), and behavior, as well as increased mortality (fish, sea urchins, krill; reviewed in Fabry et al. 2008, Pörtner 2008). The investigation of effects of ocean acidification on marine organisms is fairly recent, and only a limited number of studies address them (reviewed in Raven et al. 2005, Fabry et al. 2008, Doney et al. 2009). Even closely related species may respond differently to acidification: For example, sister species of sea urchins differ in their early life-history-stage responses to ocean acidification (reviewed in Dupont et al. 2010). This example highlights the infancy of the research in this field and our inability to generalize physiological effects. The diversity of species-specific positive and negative physiological effects in response to ocean acidification has been documented in recent reviews (table 1).

Carbon dioxide is naturally in the air: plants need it to grow, and animals exhale it when they breathe. But, thanks to people burning fuels, there is now more carbon dioxide in the atmosphere than anytime in the past 15 million years. Most of this CO2 collects in the atmosphere and, because it absorbs heat from the sun, creates a blanket around the planet, warming its temperature. But some 30 percent of this CO2 dissolves into seawater, where it doesn't remain as floating CO2 molecules. A series of chemical changes break down the CO2 molecules and recombine them with others.

When water (H2O) and CO2 mix, they combine to form carbonic acid (H2CO3). Carbonic acid is weak compared to some of the well-known acids that break down solids, such as hydrochloric acid (the main ingredient in gastric acid, which digests food in your stomach) and sulfuric acid (the main ingredient in car batteries, which can burn your skin with just a drop). The weaker carbonic acid may not act as quickly, but it works the same way as all acids: it releases hydrogen ions (H+), which bond with other molecules in the area.

Seawater that has more hydrogen ions is more acidic by definition, and it also has a lower pH. In fact, the definitions of acidification terms—acidity, H+, pH —are interlinked: acidity describes how many H+ ions are in a solution; an acid is a substance that releases H+ ions; and pH is the scale used to measure the concentration of H+ ions.

The lower the pH, the more acidic the solution. The pH scale goes from extremely basic at 14 (lye has a pH of 13) to extremely acidic at 1 (lemon juice has a pH of 2), with a pH of 7 being neutral (neither acidic or basic). The ocean itself is not actually acidic in the sense of having a pH less than 7, and it won’t become acidic even with all the CO2 that is dissolving into the ocean. But the changes in the direction of increasing acidity are still dramatic.

So far, ocean pH has dropped from 8.2 to 8.1 since the industrial revolution, and is expected by fall another 0.3 to 0.4 pH units by the end of the century. A drop in pH of 0.1 might not seem like a lot, but the pH scale, like the Richter scale for measuring earthquakes, is logarithmic. For example, pH 4 is ten times more acidic than pH 5 and 100 times (10 times 10) more acidic than pH 6. If we continue to add carbon dioxide at current rates, seawater pH may drop another 120 percent by the end of this century, to 7.8 or 7.7, creating an ocean more acidic than any seen for the past 20 million years or more.

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