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OA not OK part 1

Posted on 1 July 2011 by Doug Mackie

This post is number 1 in a series about ocean acidification. Other posts: Introduction , 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12,13, 14, 15, 16, 17, 18,Summary 1 of 2, Summary 2 of 2.

Ocean acidification, OA, is very much the other CO2 problem. Ocean acidification is the process of ocean pH decreasing (i.e. becoming more acidic) due to absorption of fossil fuel CO2 from the atmosphere. Another effect of ocean acidification is to reduce the amount of carbonate that is available to marine organisms, such as shellfish, for making their calcium carbonate shells.

The purpose of this post, the first in a twice-weekly series of posts about Ocean Acidification, is to introduce a single chemical equation. We don't want to lapse into hyperbole but the equation is the E = mc2 of ocean acidification. Having said that, knowing E=mc2 does not confer knowledge about element formation during a supernova. Likewise, this equation will only be the start of our learning:

Equation 1 with words

We need a few words about chemical equations before we discuss what Equation 1 tells us. On the left side of the arrow is the calcium ion Ca2+ (an ion is an atom or molecule that has an electrical charge, in this case 2+ meaning calcium has lost 2 negatively charged electrons to establish a +2 charge). Also on the left are 2 bicarbonate ions (HCO3-, with the negative charge meaning it has gained an extra electron). The arrow indicates that the ions on the left of the arrow react to give the molecules on the right of the arrow: calcium carbonate (CaCO3), carbon dioxide (CO2), and water (H2O).

This equation describes the formation of calcium carbonate (i.e. shells) from calcium ions and bicarbonate ions. It shows that making 1 molecule of CaCO3 from a calcium ion requires 2 molecules of bicarbonate (HCO3) and releases 1 molecule of H2O and 1 molecule of CO2. Yes, you read that right: The formation of calcium carbonate shells is a source of CO2, not a sink for CO2.

textbox 1

A basic principle is that chemical equations must be balanced. That is, they have the same number and types of atoms on both sides. Counting up we see on both the left and the right are 1 calcium (Ca), 2 hydrogen (H), 2 carbon (C), and 6 oxygen (O) atoms.

However, not all balanced chemical equations are valid chemical equations. The trick of chemistry (Oh! there's that word again) is in knowing if a particular balanced equation is valid. For example, a simple balanced equation can be written for the melting of ice at room temperature and sea level pressure:

Equation 2

But we can also write a balanced equation for the reverse reaction:

Equation 3

The problem is that, despite being able to write a balanced equation, liquid water does not spontaneously freeze at room temperature, so the second equation must be wrong. Actually what is wrong is the direction of the arrow. Reverse the arrow and Equation 2 becomes Equation 3.

Next post: the water example is familiar and we know the right answer from personal experience. But can we predict outcomes without doing the experiment? Yes – if we use thermodynamics.

Written by Doug Mackie, Christina McGraw , and Keith Hunter . This post is number 1 in a series about ocean acidification. Other posts: Introduction , 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, Summary 1 of 2, Summary 2 of 2.




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Comments 1 to 31:

  1. "Reverse the arrow and Equation 3 becomes Equation 2." Do you mean #2 becomes #3?
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  2. Best to expand the title to either: "OA (Ocean Acidification) is not OK". or "Ocean Aciditfication (OA)is not OK"
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  3. Something seems fishy here... I think the chemistry prior to Eq. 1 (ie, the formation of aqueous carbonic acid from dissolved CO2) could be written like: 2(H2O) + 2(CO2)_g <-> 2(H2CO3)_{aq} <-> 2(H+) + 2(HCO3)^{-}_{aq} Then 2 bicarbonates react with Ca ions to produce shells and returning only 1 CO2 to the atmosphere (the gas phase). Hence, if two CO2s go in, it seems like shell formation is a net carbon sink. Am I missing something?
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  4. This is a bit simplistic. You really need to introduce the concept of chemical equilibrium: Ca2+ + 2 HCO3- CaCO3 + CO2 + H2O followed by Le Chateliers principle: If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the imposed change and a new equilibrium is established. So that adding more CO2 pushes the reaction to the left - thereby dissolving CaCO3, ie: shells.
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  5. I agree, mb The whole bicarbonate equilibrium needs to be considered. Apply Le Chatelier and it's obvious that removal of bicarbonate will shift the equilibrium you've illustrated to the right, resulting in absorption of gaseous carbon dioxide from the atmosphere.
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  6. Actually, your equation (1) is most likely wrong: carbonic acid is the likely product, and then you are into the full set of carbonate equilibria: Ca2+ + 2 HCO3- CaCO3 + H2CO3     (1) CO2 + H2O H2CO3      (2) H2CO3 H+ + HCO3-      (3) HCO3- H+ + CO3     (4) Adding CO2 in eq (2) makes more H2CO3, which produces more H+ in (3) and (4), and drives the equilibrium in (1) to the left, thereby dissolving CaCO3. ps: mb et al, the markups for superscript and subscript (which makes equations much easier to read) are: superscript subscript
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  7. ps: try again: superscript subscript
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  8. Weird - it displays correctly in the preview, but then dies when posted :-/ pps: try yet again: superscript subscript
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    [DB] You were doing it right, but the WYSIWYG editor conjugates it.

    <sup> superscript </sup>
    <sub> subscript </sub>

  9. We are gratified that some find this simplistic. However, the quality of comments at this blog and elsewhere suggests that there are many who do not find this chemistry simple at all. @mb: Yes you are missing something. You are missing the point of this post which is that just because you can write a balanced equation does not mean it is a correct equation. Your set of equations is: What do you think happens to the 2H+ ? Though we invoke it later ourselves, Le Chatelier's principle is no longer taught as such in most chemistry courses. Instead it is better to compare the equilibrium constant (K) with the reaction quotient (Q).
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  10. Doug, I still can't see justification for your claim that the fromation of calcium carbonate is a source of carbon dioxide, not a sink. Could you please explain it to me again, including the formation of the bicarbonate ions in your explanation?
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  11. This one is old 1972, but if the experimental method described is sound the science is still good regardless of the date the experimentation was done. According to the author, sea water is a much more complex beast than the equations posted above. The other ions in water (such as Magnesium and Sulfate among others) influence the pH buffering system of seawater. Buffering action of sea water. A point made from author's perspective: "Control of pH in this fashion is by means of a pH-stat rather than through a buffering system. The capacity of the system is huge; at 348 ppm CO2 in the gas phase, a concentration reasonably near the normal CO2 content of the atmosphere, the extreme difference in ionic compositions produced a difference of only 0.25 pH units. Over the extreme range of CO2 concentrations, O-696 ppm, the largest pH difference was 1.4 pH units. For scawatcr with normal ionic ratios, a doubling of the CO2 content of the air would lead to a change of only 0.30 pH units an amount about cqual to the normal range of variation of seawater pH." The purpose of this post is to make sure not to miss the reality that the processes going on with seawater are more complex than a few equations.
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  12. Am I the only one thinking that criticising part 1 out of 18 of a series for being simplistic and missing things out is perhaps a bit pre-emptive?
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  13. Heraclitus @12 My post was not a criticsm of a long series of post. Just wanting to make sure they do not leave this out in future discussions.
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    [DB] Your concern is noted, but misplaced and premature.  Let us sit back and let the experts first share their knowledge before we students critique it. 

    That would be the scientific thing to do.

  14. Doug Mackie @9 I'm with mb in being a bit puzzled about the formation of CaCO3 being described as a source of CO2 rather than a sink. That statement may be true in relation to equation 1 in isolation, but if the precipitation of CaCO3 removes carbon atoms from the water, where did those carbon atoms ultimately come from? Isn't the main source the CO2 absorbed from the atmosphere? I am looking forward to this series so perhaps we will read the answer down the track.
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  15. Doug Mackie @9 One of the skeptic arguments Sks seeks to rebut is "Ocean Acidifcation Isn't Serious". On 19 March 2011, Sks posted the intermediate rebuttal under the title "Examining the Impacts of Ocean Acidifcation". This article, which I authored, describes the ocean chemistry in a set of three equations, as well as in additional notes in comment 44. While I appreciate there is more chemistry to come in this series, and different ways to write such equations, I draw your attention to these earlier articles in the interests of overall consistency of presentation of the science on this site. The material from your series will not only form the basis of a booklet, but will also be welcome input to an advanced version of the rebuttal.
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  16. DB @ 13 Will do.
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  17. 17 - good link. Some more voice
    Modeling demonstrates that if CO2 continues to be released on current trends, ocean average pH will reach 7.8 by the end of this century, corresponding to 0.5 units below the pre-industrial level, a pH level that has not been experienced for several millions of years (1). A change of 0.5 units might not sound as a very big change, but the pH scale is logaritmic meaning that such achange is equivalent to a three fold increase in H+ concentration. All this is happening at a speed 100 times greater than has ever been observed during the geological past. Several marine species, communities and ecosystems might not have the time to acclimate or adapt to these fast changes in ocean chemistry.
    I'm looking forward to the rest of this series... having been thrown out of chemistry at a very young age!
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  18. I have a copy of HSC 6.15. (A thermodynamics program) It can produce some very nice graphs of equilibrium concentrations, and effects of changing the amounts of a given component on the concentration of equlibrium species. Happy to help if requested.
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  19. HSC 6.15? How about open source? I'm pretty sure Dr Mackie knows of various sources for the original HCO3- in Eq 1. The point of the equation is what happens in the water column: the reaction increases dissolved CO2. Chemware @6, before saying Eq 1 is wrong think about the equilibrium constants for all the steps including CO2 exchange with air. The author probably teaches the details in Chem 1. Meanwhile think on the point of the equation.
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  20. Pete, Open source has the problem of what data is available, especially without scrounging. HSC has over 20000 compound and elements in its database, which is extensible. I'm sure however that open source is just fine for this case. This is rather trivial stuff. The hard part is the air/water volume assumptions you make, if you're really trying to model fluxes... and then do you partition water phases into pressure zones? My copy at this point is legacy from past employment. I keep it some statistics and graphics packages around for nostalgia's sake. I rarely use them anymore in my current roles.
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  21. @Doug_Mackie @9 - I'm not questioning the acidification of the ocean, I'm questioning the CO2 source/sink part. As you nicely show in the final line of the orange-boxed eqns., one CO2 goes in and you get sea shells plus acid (H+)... there's no returning of CO2 to the atmosphere on the RHS of the orange box rxn set. Hence I still think the original [Eq. 1] is misleading, at best, and incorrect, at worst. To answer your question ("what do you think happens to the 2 H+?"): Hydronium in the ocean can probably react with lots of stuff. But my understanding is that there's been an observed increase in [H+] (square brackets mean "concentration" for the non-chemists out there) in the ocean. Hence, I would conclude that an increase in [CO2]{g} (ie, atmospheric carbon dioxide) is in fact driving the formation of acid (actually, H3O+). Even in the theoretically closed system of the orange box with water, CO2, & calcium ions, we must conclude that an increase in [CO2]{g} on the LHS will drive the equilibrium to the RHS. Thus, CO2{g} is being taken from the atmosphere, not to be returned (or only returned in small quantities... chemists would draw that conclusion as a big arrow pointing to the right and a small arrow pointing left). More generally, (and as has been pointed out) this discussion is somewhat without context if we don't discuss rate constants and concentrations. That said, there's not reason to believe OA is, prima facia, a carbon source. Instead, OA is a result of increased [CO2], driving the last line of the orange box to the RHS (to the acid-producing side) a-la LeChat's (ie, big arrow pointing to the RHS). (Incidentally, the orange-boxed equn. in comment 9 is missing a 2 on the first line in front of CO2).
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  22. Also, I should add that much of my previous post assumes an unlimited supply of calcium ions. So.... crush up a bunch of Tums and put it in the ocean to eat up CO2! :P
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  23. mb... you need to keep in mind that solubility of common calcium species in water is far, far less than say, sodium Here's a quick table of solubility product constant Ksp of common species: Calcium carbonate CaCO3 3.8 x 10-9 Calcium flouride CaF2 3.4 x 10-11 Calcium hydroxide Ca(OH)2 7.9 x 10-6 Calcium oxalate CaC2O4 1.5 x 10-8 Calcium phosphate Ca3(PO4)2 1.0 x 10-26 Calcium sulfate CaSO4 2.4 x 10-5 Precipitating out CaCarbonate doesn't make it immediately available to to the species that need it. Take a look at this wikipedia article on calcium carbonate= WikiPedia Article it's relatively good except that it leaves out effect of water temperatuer on the solubility product constant- calcium carbonate is unusual in that Ksp decreases with increasing temperature.
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  24. We know CO2-enriched sea water causes ‘seashells’ (CaCO3) to erode. link . Therefore CaCO3 + CO2 in sea water --> something that isn’t CaCO3 [solid] This suggests the possibility (reversing the arrow) that growing ‘seashells’ (CaCO3) causes the release of CO2. This may not be vigorous chemistry, but the field work supports the blog's conclusion. [Geologists approach things with a geopick or a sledgehammer.]
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  25. Alan@15 The main source of carbon for calcification in the ocean is the bicarbonate that is already there. Figure 7.3 from the IPCC 4th Assessment Report (2007) reminds us that the oceans contain more than 50x as much carbon as the atmosphere. Later posts will relate equation 1 with the shorthand way to describe the formation of calcium carbonate ( Ca2+ + CO3 --> CaCO3). This simplified equation ignores the fact that actual calcification involves the consumption of bicarbonate and release of CO2. In order to for everyone to follow the chemical and physical processes behind ocean acidification (the goal of this series), we need to start with these individual reactions. By the end of the series, you'll find that the equations will tie in nicely with your posts. mb@21: Yes, [H+] has increased. BUT it has increased at less than purely stoichiometric calculations would suggest. We explain what that means and why it is so in future posts. Others: This is post 1 of 18. We will address your concerns. We all deride blog science. Blog science is what happens when explanations are rushed. Real science takes time to explain. As we said in the introduction, the chemistry is deceptively complex but have patience and we will get there.
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  26. I really appreciate that you're providing a primer on the chemistry of ocean acidification. It is sorely misunderstood by most people, and I'm glad to have this resource to direct folks to.
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  27. I happened across the following article posted on redOrbit. Everyone reading this thread will find it interesting. Climate Change Could Turn Oxygen-Free Seas From Blessing To Curse For Zooplankton
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  28. Norman@11, While I agree with you that the chemistry is more complex than just those equations involved. Your article is very old. 1972. It has now been seen that at 8.00 pH aragonite forming organisms are no longer able to make their shells. Between 1751 and 1994 surface ocean pH is estimated to have decreased from approximately 8.25 to 8.14 wiki/Ocean_acidification The attitude that 0.25 pH units is small is now totally out of date. It is considered that 0.09 pH change is large for the ocean! Only half that change you speak of as small will wipe out aragonite life forms from the current state of the oceans. (except for refuges)
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  29. Paul W@28: We will discuss this in post 16 (long time to wait I know) but briefly: Calcification is better described as a function of carbonate concentration (or saturation state, omega) than a function of pH. In surface water for constant alkalinity typical of surface seawater (more on this in a later post) a pH of 8 implies about 190 umol carbonate per kg of seawater. At pH 8.14, there is about 150 umol carbonate /kg. So you are correct - seemingly small changes pH can have noticeable differences in carbonate concentrations. But, calcification can still happen at these levels – see later posts
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  30. second summary post

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  31. Could someone check out this interactive on sea chemisty and tell me if its incorrect?


    The interactive on the right, thank you.

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