Some organisms may tolerate a certain amount of change, but thinner shells will make others more vulnerable to damage or predators. Some organisms might also tolerate acidification of internal fluids to a point, yet even so many will expend more energy to maintain their optimal acid-base balance or will struggle to supply their body with oxygen and to sustain cellular functions vital to life. The extra expense of coping with acidification may make them more prone to dying. These stresses will be particularly severe for deep-sea animals, which have adapted to an extremely stable environment. And even if animals survive, the stresses will sap energy they would otherwise use for growth and reproduction.
We would probably see the effects of ocean acidification first in animal groups that have finely tuned environmental ranges, particularly those already “living on the edge” such as coral reefs, which have already suffered widespread bleaching and death from warming ocean temperatures. Less appreciated are effects on massive communities of tiny animals that live in the ocean’s midlevels. These creatures migrate en masse to the surface layer at night to feed yet sink to deep water during the daytime to avoid predators. In so doing, they form a critical link between the warm, oxygenated surface layer and the cold, oxygen-depleted waters of the deep, as well as a critical link in the oceanwide food chain.
Increased acidity and expanding zones of low oxygen in some regions may force these midwater organisms into shallower waters where they would be more exposed to predators. And if, as expected, the zones of low oxygen expand and intensify, many of these migrators could die. Together these effects could slice through this daily, migratory lifeline between shallow and deep waters—an outcome that could impact society’s ocean fisheries.
How well marine life can adapt to rapid acidification remains an open question, but there is real reason for concern. Ocean life has weathered large environmental perturbations during the earth’s history, just barely; some 250 million years ago massive volcanism is thought to have caused ocean acidification and other factors that left 90 percent of marine species dead.
Although man-made climate change will be much milder, strong and immediate action to stabilize CO2 levels is essential to minimize our disruption of ocean chemistry and ecosystems. We can no longer deny our role in global climate change. Now is the time for serious discussion among science, business and political leaders about ways to minimize our impact on our air and water, to set limits on the effects of our fossil-fuel use, and to plan how to adapt to coming change.
Note: This article was originally printed with the title, "The Other CO2 Problem".
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19 Comments
Add CommentThis claim of ocean absorption of CO2 causing reduced levels of carbonate ions has to be news to every trained chemist.
Reply | Report Abuse | Link to thisHow can adding CO2 to seawater possibly cause a reduction in carbonate ions? When you add CO2 to water some of it reacts with the water to form acidic protons and carbonate ions. The reaction is:
CO2 + H20 = 2H+ + CO3--
The equilibrium constant equation is:
[CO2] = k x [H+] x [H+] x [CO3--] where k is the dissociation or equilibrium constant and the bracketed terms represent the concentrations of those entities.
Adding CO2 must increase the concentration of carbonate ions.
The only way carbonate ion concentration can be reduced is if acid is added from some other source or if carbonate is removed by some other process.
This is basic solution chemistry. Why give credence to unexplained chemical nonsense?
You've neglected the amphiprotic nature of CO3-. ie, HCO3- + H2O --> CO32- + H3O+ and HCO3- + H2O > H2CO3 + OH-. The question is what are those equilibriums, and what are their corresponding strengths compared to the atmospheric pressure of co2.
Reply | Report Abuse | Link to thisOur bodies expose our blood to the same co2. Our blood has the exact same salinity. Anyone notice the mass retardation? So you think you might have a little acid reflux? Math scores dropping lately? Bankers doing dumb stuff? Having a little antibiotic resistant bacteria? You humans have a nice day.
Reply | Report Abuse | Link to thisPraxis, the bicarbonate equilibrium doesn't make any difference to the equilibrium between CO2 and carbonate ion. It is already accounted for in the carbonate equilibrium constant.
Reply | Report Abuse | Link to thisPut simply, there is no way that adding CO2 to water reduces carbonate concentration. It may well reduce Calcium concentrations and that may affect shell growth for all I know but the argument in this article about carbonate concentrations is just utter nonsense.
The principle problem is that dissolved CO2 is dissolving carbonate in the form of shells or plankton skeletal structures. The exact chemical equilibria involved may not be known, but the damage is there.
Reply | Report Abuse | Link to thisThat is totally absurd, eco-steve. More CO2 implies greater carbonate concentrations which will inhibit rather than encourage shells from dissolving. You will have to look deeper for the cause, assuming that the damage has more reality than the present ludicrous explanation.
Reply | Report Abuse | Link to thisHow in the world can SciAm publish an article that doesn't even make sense? This is ridicuous...
Reply | Report Abuse | Link to thisok, i feel a little silly, but I rescind my previous comment, sort of. The author does not do a good job explaining what he means by 'the carbonate concentration will decrease'. It would have been accurate to state that the carbonate ion concentration relative to total dissolved carbon decreases with increasing CO2 absorption.
Reply | Report Abuse | Link to thisYou were right the first time, enochdames. What mechanism or evidence is there that carbonate ion concentration relative to anything decreases with increasing CO2 absorption?
Reply | Report Abuse | Link to thisThe equilibrium constant I quoted establishes a linear relationship between CO2 and carbonate ion concentrations.
This article is chemical nonsense, pure and simple.
the equation you wrote is valid, but it is one of many in the aqueous carbonate system. take a look at this link (http://lawr.ucdavis.edu/classes/ssc102/Section5.pdf) and check out the figure on the 5th page. In short, the carbonate ion is a conjugate base of the bicarbonate ion, and increasing acidity decreases its relative concentration
Reply | Report Abuse | Link to thisThat's true, enochdames, but since both carbonate and bicarbonate conentrations increase with increasing CO2 how is that relevant to shell formation issues?
Reply | Report Abuse | Link to this(I acknowledge my previous post was hasty in that the linearity only exists if acidity is unchanged.)
Reply to Alanw: No doubt it may not seem logical, but increasing dissolved CO2 in seawater preferentially dissolves the shells of marine organisms. This is the conclusion after many tests. Experiment is always more valid that theoretical guesswork....
Reply | Report Abuse | Link to thiseco-steve, fine if that is true in which case they should present the data and let it speak for itself rather than advance a pseudo-explanation that is scientifically absurd.
Reply | Report Abuse | Link to thisWhy in this article would you even put this sentance "some 250 million years ago massive volcanism is thought to have caused ocean acidification and other factors that left 90 percent of marine species dead."
Reply | Report Abuse | Link to thisAs the next paragraph admits the process that brought about 90% extinction rates is on a scale and of a nature completely unrelated to the CO2 acidification process that is under discussion in the rest of the article.
The general need by environmental activists to exaggerate what is going on and get sensationalist about things suggests to me a weakness in the real supportive data. It's about time we saw less hype and more science.
Seems to me the larger concern here is the effect rising ocean acidity will have not on larger shellfish or snails, but on diatoms. Diatoms produce most of the oxygen in the atmosphere. Collapse of their populations could have drastic effects on all oxygen-breathing life on the planet.
Reply | Report Abuse | Link to thisDespite your thoughts, CO2 can't be a good source to rising the acidity of the ocean...
Reply | Report Abuse | Link to thisDespite your thoughts, CO2 can't be a good source to rising the acidity of the ocean...
Reply | Report Abuse | Link to thisThis has to be about the most "unscientific" drivel I have read in years.
Reply | Report Abuse | Link to thisFACT 1: The solubility of a gas - any gas - in a liquid - any liquid, is inversely proportional to the temperature of the liquid, (assuming more or less constant atmospheric pressure). This is one of the so-called "Gas Laws".
Therefore the "amount" of CO2 (regardless of source) dissolved in the ocean is DIRECTLY dependent on the TEMPERATURE of the water, and on NO other factor .
FACT 2: If indeed the global temperatures are rising (including the oceans) - as some would have us believe, then the total amount of dissolved CO2 will - MUST - DECREASE, not INCREASE, and atmospheric levels of CO2 WILL increase.
Conversely, if global temperatures are decreasing, as some argue, the levels of dissolved CO2 WILL increase, and levels of atmospheric CO2 WILL decrease.
Since there would have been a vastly higher level of dissolved CO2 in the oceans during the last Ice Age than now, and since every species alive now in the oceans survived that, preposterous arguments about "man-made ocean acidification" causing damage can be seen for what they are - asinine, entirely non-scientific attempts by the clueless to jump on the AGW gravy train.
There seems to be some confusion among people regarding the effect of increasing CO2 concentration in the atmosphere on the oceans.
Reply | Report Abuse | Link to thisThe first effect to look at is the direct effect of increasing the atmosphere's CO2 concentration on the ocean. The amount that dissolves in a liquid is described by Henry's Law.
"At constant temperature the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid"
This is basically saying that the higher the concentration of the gas the more will dissolve in the liquid and conversely the more gas dissolved in the liquid the less gas the liquid will absorb. If the liquid has a very high level of gas and its partial pressure is higher than the surrounding atmosphere then the liquid will release that gas.
How temperature effects solubility is actually quite complicated but can be simplified; At around room temperature, increasing the temperature reduces the solubility of the gas. For sea water it takes an increase of 16K to half the amount of gas the liquid can absorb. As CO2 has increased by 70% over the last 200 years it would take an increase of 10K to get the amount of CO2 in the seas back to how it was. So far, due to the seas immense weight and therefore heat capacity the seas have increased in temperature by far less than 1K. Therefore as has been measured over the past decades we will expect CO2 levels in the sea to increase.
As CO2 increases so the pH decreases. As more CO2 is added more bicarbonate and carbonate is added to the oceans. As more CO2 is added bicarbonate becomes more favoured and carbonate less so. As sea shells are made of carbonate increasing the CO2 levels causes the shells to dissolve more readily.