Ocean Acidification: Can CO2 Release & pH Decrease Simultaneously?

[cesiumfrog]

How can the ocean simultaneously release CO2 and decrease in pH?

Currently the ocean is acidifying, as it absorbs about a third of the fossil-carbon dioxide that we emit, which then in part assumes the form of carbonic acid. But in the future, if the increasing atmospheric greenhouse effect continues to also warm the ocean enough, we expect this absorption will be reversed and vast quantities of CO2 will distil out from the ocean.

Nonetheless, apparently we do not expect warming to cause any reverse to the acidification. (I asked one paleoceanographer/marine-chemist, and heard there is no contradiction for water to be simultaneously decreasing in pH and liberating CO2.) But naively, if warmed water begins losing carbon, then shouldn't the concentration of carbonic acid fall (and hence the pH start to rise back again)?

Edit: The topic of this thread is not "global warming or climate change". It is purely an ocean-chemistry question. Regardless of what is actually happening to our ocean (or rather, regardless of what external factors may be controlling the temperature of and the partial pressure of CO2 above a hypothetical test-ocean) the question is simply whether in principle such an ocean hypothetically could ever be driven (by adjusting those two parameters) to release CO2 while simultaneously to decrease in pH, and how exactly? (So this is what self-censorship is like..)

 

[Xnn]

Currently, the ocean is absorbing CO2 from the atmosphere and its pH is decreasing.
As the ocean warms, it ability to absorb CO2 will become diminished.
Eventually, if it warmed enough, it could begin to release CO2.
However, that'd be a long time in the future and the pH would have become stable at a much lower value.

So, we are not currently witnessing a simultaneous release of CO2 and decreasing pH.




http://www.ipcc.ch/pdf/assessment-report/ar4/wg1/ar4-wg1-ts.pdf

See page 48:

The uptake of anthropogenic carbon since 1750 has
led to the ocean becoming more acidic, with an average
decrease in surface pH of 0.1 units. Uptake of CO2 by
the ocean changes its chemical equilibrium. Dissolved
CO2 forms a weak acid, so as dissolved CO2 increases,
pH decreases (i.e., the ocean becomes more acidic).

The overall pH change is computed from estimates of
anthropogenic carbon uptake and simple ocean models.
Direct observations of pH at available stations for the
last 20 years also show trends of decreasing pH, at a rate
of about 0.02 pH units per decade. Decreasing ocean
pH decreases the depth below which calcium carbonate
dissolves and increases the volume of the ocean that is
undersaturated with respect to the minerals aragonite
(a meta-stable form of calcium carbonate) and calcite,
which are used by marine organisms to build their
shells. Decreasing surface ocean pH and rising surface
temperatures also act to reduce the ocean buffer capacity
for CO2 and the rate at which the ocean can take up excess
atmospheric CO2.

 

[sylas]

[strike]Unfortunately, this topic is one physicsforums feels unable to cope with.[/strike] I stand corrected; my apologies. Thanks, Evo.

As the ocean warms, it ability to absorb CO2 will become diminished.
Eventually, if it warmed enough, it could begin to release CO2.

However, just for the record, there is no prospect of the ocean releasing dissolved CO2 unless atmospheric levels are much lower. The only prospect is a reduced capacity to absorb CO2.

 

[Studiot]

Unfortunately, this topic is one physicsforums feels unable to cope with.

Excuse me?

 

[Evo]

Unfortunately, this topic is one physicsforums feels unable to cope with.

The question about the effect of CO2 on ocean acidity, as posed by the OP, is perfectly fine.

 

[Studiot]

The problem I see is one of the sheer size of the ocean.

To this day there are soundings on the Admiralty charts that were taken by Captain Cook.
That part of the ocean has not been resounded let alone had its chemical composition measured.

I just think our sample sizes are way too small to make such claims about the composition of the ocean, with any confidence, let alone changes in composition of several orders of magnitude finer graining.

 

[sylas]

The problem I see is one of the sheer size of the ocean.

To this day there are soundings on the Admiralty charts that were taken by Captain Cook.
That part of the ocean has not been resounded let alone had its chemical composition measured.

I just think our sample sizes are way too small to make such claims about the composition of the ocean, with any confidence, let alone changes in composition of several orders of magnitude finer graining.

The measurements are plenty good enough to identify the falling pH levels. The chemistry of how this occurs is also well known. The increase in atmospheric CO2 is very well measured indeed, and the carbon cycle which exchanges carbon between ocean and atmosphere means that the ocean is currently increasing in total carbon content as well. This inevitably leads to a falling pH.

The carbon cycle also means that there's an equilibrium maintained between the oceans and the atmosphere, and given the rapidly rising atmospheric carbon levels, there is a large excess in atmospheric carbon from the equilibrium state. That's why the ocean is removing carbon from the atmosphere, and also why there is no credible prospect of that reversing. The only issue is whether that rate of ocean take up might slow down as levels in the ocean continue to increase. So far it seems to be holding pretty steady.

Hence as atmospheric carbon increases, ocean carbon increases, and ocean pH falls. This much is straightforward.

Of course, that's a pretty broad description, and scientists want to measure the rates of change more precisely, and there's a lot of work and data on this as well.

See, for example:

• Wootton, J.T. et al (2008) http://210.193.216.98/cps/rde/papp/...http://www.pnas.org/content/105/48/18848.full, in PNAS, Vol 105, no. 48, pp 18848-18853.

Cheers -- sylas

 

[Studiot]

The measurements are plenty good enough to identify the falling pH levels.

Would these be the same measurements that 15/20 years ago lead to the belief/claim that life could not exist at the depths and temperatures of the thermal plumes at the ocean bed?

 

[sylas]

Would these be the same measurements that 15/20 years ago lead to the belief/claim that life could not exist at the depths and temperatures of the thermal plumes at the ocean bed?

I guess that is meant to be a rhetorical question, since the answer is so obviously of course not. More importantly, it wasn't measurements that were in question, but finding examples of living things that were able to tolerate those conditions.

Back to the actual question of the thread... measurements show without any doubt that carbon levels in the atmosphere and in the ocean are increasing, and that pH is decreasing as well. You can't have one without the other; the chemistry means more carbon in the ocean corresponds to a lower pH.

Cheers -- sylas

PS. And on the side issue of the fascinating history of hydrothermal plumes. The plumes were discovered in 1977, and life in those plumes was discovered at the same time.

 

[Gokul43201]

Would these be the same measurements that 15/20 years ago lead to the belief/claim that life could not exist at the depths and temperatures of the thermal plumes at the ocean bed?

1. Beliefs/claims are relevant to any discussion only if there are printed in peer reviewed publications or standard textbooks. All other claims are irrelevant.

2. Any assertions about claims need to be accompanied by citations to the papers/texts that they appear in. All other assertions are irrelevant.

 

[Evo]

Would these be the same measurements that 15/20 years ago lead to the belief/claim that life could not exist at the depths and temperatures of the thermal plumes at the ocean bed?

Here is information about life at thermal vents that you might find interesting.

http://www.waterencyclopedia.com/La-Mi/Life-in-Extreme-Water-Environments.html

 

[Studiot]

Thank you for the link, Evo.

I was not thinking of micro-organisms. I was really thinking about the BBC videos showing the 'fish' swimming near these spots.

But you are right, any life is still life.

I have just been reading Sylas' excellent blog on debate and communication; I am just trying to make the point how little information we really know about the deeps.
A peer review of 3, 30 or 3000 measurements in an ocean the size of the Pacific is pointless.

Can anyone say how many measurements we take of the composition of the bottom 50%of the Pacific water per decade?
Divide this by the volume of that water and you will have a figure of measurements per cubic mile or whatever. The figure will not be large enough in my opinion to statistically state that the composition of the Pacific is changing in any direction.

It may be (probably is) that some locations are sampled more frequently so such a statement could be made confidently about such a location. But it would be a statistical fallacy to use this to represent the ocean as a whole.

 

[sylas]

Thank you for the link, Evo.

I was not thinking of micro-organisms. I was really thinking about the BBC videos showing the 'fish' swimming near these spots.

But you are right, any life is still life.

I have just been reading Sylas' excellent blog on debate and communication; I am just trying to make the point how little information we really know about the deeps.
A peer review of 3, 30 or 3000 measurements in an ocean the size of the Pacific is pointless.

Can anyone say how many measurements we take of the composition of the bottom 50%of the Pacific water per decade?
Divide this by the volume of that water and you will have a figure of measurements per cubic mile or whatever. The figure will not be large enough in my opinion to statistically state that the composition of the Pacific is changing in any direction.

It may be (probably is) that some locations are sampled more frequently so such a statement could be made confidently about such a location. But it would be a statistical fallacy to use this to represent the ocean as a whole.

Thanks for the comment on my blog article! It's something I've often thought about, and how discussions like this can be productive even with disagreements and even when they don't seem to be actually resolved.

I expect the pH of the deep ocean is uncertain; but what is quite certain is the following.

(A) The ocean is soaking up one heck of a lot of carbon, at the rate of about 2 Gton per year. This is absorbed mainly at the surface, from the atmosphere, and then circulates through the rest of the water column; although the rate at which carbon is accumulating at different depths is not completely clear. Roughly speaking, new carbon is being added to the carbon cycle at about 8 Gt per year from geological sources (fossil fuels). Almost half remains in the atmosphere (giving a well measured increase) and the rest is about half each into oceanic and terrestrial carbon sinks.

Details of how carbon is increasing in the reservoirs of the carbon cycle is considered in (for example)

• Knorr, W. (2009), http://dx.doi.org/10.1029/2009GL040613, in Geophys. Res. Lett., 36, L21710, doi:10.1029/2009GL040613.
• Bake, D.F. (2007), http://www.sciencemag.org/cgi/content/summary/316/5832/1708, in Science vol 316, pp 1708-1709, doi:10.1126/science.1144863

(B) Adding carbon in the ocean is chemically fairly straightforward; it ends up as carbonic acid, and reduced the pH of the ocean. This is measured mainly in the upper ocean, where we also can see its effects most clearly, and where the impacts are likely to be greatest.

You don't need to know all the details of all the ocean to know that it is absorbing one heck of a lot of carbon, and that this lowers the pH. Details of how it is distributed are a perfectly good question, but the answers to the original thread question are still pretty straightforward. Specifically, there is no prospect of changing ocean temperatures reversing the currently measured trends of ocean water absorbing carbon and reducing in pH.

Cheers -- sylas

 

[Studiot]

Sylas I don't disagree with your assesment of the atmousphere - ocean interaction.

But the ocean has a bottom as well as a top. And as I understand the current theories of geological activity, movement of most of the carbon in the world occurs in carbonate rock in the bottom part. I also understand that the disposition of the continents plays a bigger role than is generally recognised in the generation of (extreme) weather.

We have seen at least one Earth shattering (no pun intended) discovery in geology every decade in the second half of the 20th century which has stood conventional thinking on its head.

The outcropping of these dicoveries seems unabated as we enter the 21st century,
for instance the work on the PT bounday and the mass extinction at that time.

 

[sylas]

Sylas I don't disagree with your assesment of the atmousphere - ocean interaction.

OK... but I would have thought that was the underlying basis of the question of the thread.

But the ocean has a bottom as well as a top. And as I understand the current theories of geological activity, movement of most of the carbon in the world occurs in carbonate rock in the bottom part. I also understand that the disposition of the continents plays a bigger role than is generally recognised in the generation of (extreme) weather.

Quite so. But geological activity occurs at much too slow a rate to have much of an impact on anything in the time frames for the question being asked here.

Most of the carbon may be bound up in carbonate rock; but because this exchanges only very slowly, such rock has little impact on the carbon cycle except over extremely long time scales. There is very little actual movement of carbon in and out of rock, by comparisons with the enormous fluxes of carbon dominating the carbon cycle.

The shape of land forms certainly has an impact on weather in all kinds of ways; and this means weather patterns in the past might be radically different. But the disposition of continents changes slowly, and continents are going to remain arranged pretty much as we have them now for a long time to come. I had taken the question to be about the impact of the ocean on carbon cycles during the coming century or millennium.

We have seen at least one Earth shattering (no pun intended) discovery in geology every decade in the second half of the 20th century which has stood conventional thinking on its head.

The outcropping of these dicoveries seems unabated as we enter the 21st century,
for instance the work on the PT bounday and the mass extinction at that time.

As a general principle in science, everything we think we know might be wrong. But its unlikely for it ALL to be wrong, and the fact that we do indeed make discoveries shows that we there must be some kind of a basis for answering scientific questions such as the one in this thread.

The original question was:

How can the ocean simultaneously release CO2 and decrease in pH? ​

The short answer is... only if CO2 is not actually coming from the ocean itself, but from something else. For example, a large subterranean volcano might suddenly go off, and release a lot of carbon, some of which will remain in the water (and hence lower the pH) and some of which might be released into the atmosphere. Same for melting calthrate deposits, or any other such unexpected process.

The original post of the thread includes a statement which I believe is incorrect.

Currently the ocean is acidifying, as it absorbs about a third of the fossil-carbon dioxide that we emit, which then in part assumes the form of carbonic acid. But in the future, if the increasing atmospheric greenhouse effect continues to also warm the ocean enough, we expect this absorption will be reversed and vast quantities of CO2 will distil out from the ocean

I do not think that second sentence is true. There is no expectation that lots of CO2 will distill out of the ocean; what may happen is that the rate at which it is being absorbed will slow down. As water temperatures increase, solubility of carbon reduces. This is a likely cause of the large changes in atmospheric carbon dioxide during the ice ages. However, with the present equilibirum having so large an excess of atmospheric carbon dioxide, the natural equilibrium is for the ocean to continue to remove atmospheric carbon. Increasing temperatures might reduce the total amount that could be absorbed, but not reverse the direction of net absorption.

Reference, on changing solubility of carbon dioxide with temperature and the consequent feedback loop between temperature and atmospheric CO2:

• Martin, P.D. et. al. (2005), http://www.agu.org/pubs/crossref/2005/2003PA000914.shtml, in Paleoceanography, 20, PA2015, doi:10.1029/2003PA000914.

As matters stand, there is one heck of a lot of carbon is being absorbed by the ocean, as CO2 is dissolved at the surface and becomes carbonic acid... lowering the pH.

For the ocean to actually emit more carbon than it is absorbing, something radically new and different would have to occur -- not merely release of dissolved carbon. Something unexpected. Not merely something we haven't noticed in existed measurements. For now, the ocean is absorbing a bit over 2 Gt of carbon per year, and geological outgassing is negligible by comparison with that. The absorbed carbon reduces pH.

Cheers -- sylas

 

[Evo]

Specifically, there is no prospect of changing ocean temperatures reversing the currently measured trends of ocean water absorbing carbon and reducing in pH.

Cheers -- sylas

Actually a study was published in the journal Science January 16, 2009, stating that if the oceans warmed as predicted that they expect ocean alkalinity to increase, but this definitely crosses the boundary into predictions of climate change, but since it addresses the topic, I will post a link to the article discussing the study, but this is purely for information on the study, not to start a discussion on CC. We can defnitely discuss the effects of fish on ocean alkalinity.

http://www.sciencedaily.com/releases/2009/01/090115164607.htm

Link to abstract in Science. http://www.sciencemag.org/cgi/content/abstract/sci;323/5912/359?maxtoshow=&hits=10&RESULTFORMAT=&fulltext=ocean+acidification+fish+january+16%2C+2009&searchid=1&FIRSTINDEX=0&resourcetype=HWCIT

Also, to address studiot's question about acidity levels at various places, shallower water will tend to be more alkaline.

 

[sylas]

Actually a study was published in the journal Science January 16, 2009, stating that if the oceans warmed as predicted that they expect ocean alkalinity to increase, but this definitely crosses the boundary into predictions of climate change, but since it addresses the topic, I will post a link to the article discussing the study, but this is purely for information on the study, not to start a discussion on CC. We can defnitely discuss the effects of fish on ocean alkalinity.

I'm quite happy to stick with chemistry here. The biggest issue here is a potential confusion of alkalinity and pH. They are defined differently, in terms of different ions.

Thanks for the link! I checked it out. The reference discussed in that press release is:

• Wilson R.W. et. al. (2009) http://www.sciencemag.org/cgi/content/abstract/323/5912/359, in Science, Vol. 323. no. 5912, pp. 359 - 362, doi:10.1126/science.1157972

Abstract:

Oceanic production of calcium carbonate is conventionally attributed to marine plankton (coccolithophores and foraminifera). Here we report that marine fish produce precipitated carbonates within their intestines and excrete these at high rates. When combined with estimates of global fish biomass, this suggests that marine fish contribute 3 to 15% of total oceanic carbonate production. Fish carbonates have a higher magnesium content and solubility than traditional sources, yielding faster dissolution with depth. This may explain up to a quarter of the increase in titratable alkalinity within 1000 meters of the ocean surface, a controversial phenomenon that has puzzled oceanographers for decades. We also predict that fish carbonate production may rise in response to future environmental changes in carbon dioxide, and thus become an increasingly important component of the inorganic carbon cycle.​

As the press release seemed to indicate, this research is dealing mainly with alkalinity rather than pH. I don't see that this research conflicts with any of what has been said already in the thread. The bearing on the original questions seems to quite limited, unless I am missing something drastic.

Cheers -- sylas

 

[Evo]

I'm quite happy to stick with chemistry here. The biggest issue here is a potential confusion of alkalinity and pH. They are defined differently, in terms of different ions.

Thanks for the link! I checked it out. The reference discussed in that press release is:

• Wilson R.W. et. al. (2009) http://www.sciencemag.org/cgi/content/abstract/323/5912/359, in Science, Vol. 323. no. 5912, pp. 359 - 362, doi:10.1126/science.1157972

Abstract:

Oceanic production of calcium carbonate is conventionally attributed to marine plankton (coccolithophores and foraminifera). Here we report that marine fish produce precipitated carbonates within their intestines and excrete these at high rates. When combined with estimates of global fish biomass, this suggests that marine fish contribute 3 to 15% of total oceanic carbonate production. Fish carbonates have a higher magnesium content and solubility than traditional sources, yielding faster dissolution with depth. This may explain up to a quarter of the increase in titratable alkalinity within 1000 meters of the ocean surface, a controversial phenomenon that has puzzled oceanographers for decades. We also predict that fish carbonate production may rise in response to future environmental changes in carbon dioxide, and thus become an increasingly important component of the inorganic carbon cycle.​

As the press release seemed to indicate, this research is dealing mainly with alkalinity rather than pH. I don't see that this research conflicts with any of what has been said already in the thread. The bearing on the original questions seems to quite limited, unless I am missing something drastic.

Cheers -- sylas

pH measures both acidity and alkalinity. A low pH number indicates acidity and a high pH rating indicates alkalinity. Technically, you could say pOH for alkalinity, but the article the OP was referring to just uses a scale with high pH = alkaline, low = acid.

To get more specific, there are different ways to measure ocean water.

Seawater
The pH of seawater plays an important role in the ocean's carbon cycle and there is evidence of ongoing ocean acidification caused by carbon dioxide emissions.[19] However, pH measurement is complicated by the chemical properties of seawater, and several distinct pH scales exist in chemical oceanography.[20]

As part of its operational definition of the pH scale, the IUPAC define a series of buffer solutions across a range of pH values (often denoted with NBS or NIST designation). These solutions have a relatively low ionic strength (~0.1) compared to that of seawater (~0.7), and consequently are not recommended for use in characterising the pH of seawater since the ionic strength differences cause changes in electrode potential. To resolve this problem, an alternative series of buffers based on artificial seawater was developed.[21] This new series resolves the problem of ionic strength differences between samples and the buffers, and the new pH scale is referred to as the total scale, often denoted as pHT.

The total scale was defined using a medium containing sulfate ions. These ions experience protonation, H+ + SO42− ⇌ HSO4−, such that the total scale includes the effect of both protons (free hydrogen ions) and hydrogen sulfate ions:

[H+]T = [H+]F + [HSO4−]

An alternative scale, the free scale, often denoted pHF, omits this consideration and focuses solely on [H+]F, in principle making it a simpler representation of hydrogen ion concentration. Analytically, only [H+]T can be determined,[22] therefore, [H+]F must be estimated using the [SO42−] and the stability constant of HSO4−, KS*:

[H+]F = [H+]T − [HSO4−] = [H+]T ( 1 + [SO42−] / KS* )−1

However, it is difficult to estimate KS* in seawater, limiting the utility of the otherwise more straightforward free scale.

Another scale, known as the seawater scale, often denoted pHSWS, takes account of a further protonation relationship between hydrogen ions and fluoride ions, H+ + F− ⇌ HF. Resulting in the following expression for [H+]SWS:

[H+]SWS = [H+]F + [HSO4−] + [HF]

However, the advantage of considering this additional complexity is dependent upon the abundance of fluoride in the medium. In seawater, for instance, sulfate ions occur at much greater concentrations (> 400 times) than those of fluoride. Consequently, for most practical purposes, the difference between the total and seawater scales is very small.

The following three equations summarise the three scales of pH:

pHF = − log [H+]F

pHT = − log ( [H+]F + [HSO4−] ) = − log [H+]T

pHSWS = − log ( [H+]F + [HSO4−] + [HF] ) = − log [H+]SWS

In practical terms, the three seawater pH scales differ in their values by up to 0.12 pH units, differences that are much larger than the accuracy of pH measurements typically required, particularly in relation to the ocean's carbonate system.[20] Since it omits consideration of sulfate and fluoride ions, the free scale is significantly different from both the total and seawater scales. Because of the relative unimportance of the fluoride ion, the total and seawater scales differ only very slightly.

http://en.wikipedia.org/wiki/PH#Mathematical_Definition

 

[Xnn]

However, just for the record, there is no prospect of the ocean releasing dissolved CO2 unless atmospheric levels are much lower. The only prospect is a reduced capacity to absorb CO2.

That's a good point, but I don't claim to know if that is true for sure.

What I do know is that as water warms, its ability to hold dissolved gases, such as CO2 is diminished. So, theoretically, atmospheric CO2 levels do not need to fall for the oceans to release CO2. If the oceans were to warm enough, they could begin to release CO2 even if atmospheric CO2 levels were to continue to rise. It's all a matter of the relative response of each characteristic.

There is also the question of how the oceans will warm. Will they stagnate or not (and if so, then to what degree)?

If the oceans were to stagnate significantly, then there would be an exceptionally warm surface layer with very low solubility for CO2 atop a colder layer that could be under saturated.

 

[Borek]

Chemically I can't think of a way of decreasing amount of CO₂ in water AND decreasing pH. As Sylas pointed out, more dissolved carbon dioxide means more carbonic acid means lower pH, period. But I guess we can be talking apples and oranges, that is, paleoceanographer/marine-chemist asked by OP was probably not thinking about simple model (bucket of water and air with CO₂ above), but much more complicated system with additional sources and sinks of CO₂ (biomass, precipitation and so on).

 

[Evo]

Chemically I can't think of a way of decreasing amount of CO₂ in water AND decreasing pH. As Sylas pointed out, more dissolved carbon dioxide means more carbonic acid means lower pH, period. But I guess we can be talking apples and oranges, that is, paleoceanographer/marine-chemist asked by OP was probably not thinking about simple model (bucket of water and air with CO₂ above), but much more complicated system with additional sources and sinks of CO₂ (biomass, precipitation and so on).

Read the Science article Borek, it's in my post. It's not about decreasing CO₂.

 

[Borek]

pH measures both acidity and alkalinity.

I think that's incorrect. Two solutions with different pH can have the same alkalinity, two solutions with the same pH can have different alkalinity. There is no simple dependency between both.

Please remember that alkalinity measures (more or less) buffering capacity of the solution, that is it measures how the concentration of H⁺ changes when acids and bases are added, while pH measures just concentration of H⁺.

 

[Evo]

I think that's incorrect. Two solutions with different pH can have the same alkalinity, two solutions with the same pH can have different alkalinity. There is no simple dependency between both.

Please remember that alkalinity measures (more or less) buffering capacity of the solution, that is it measures how the concentration of H⁺ changes when acids and bases are added, while pH measures just concentration of H⁺.

Yes, and I edited my post

Technically, you could say POH for alkalinity, but the article the OP was referring to just uses a scale with high pH = alkaline, low = acidity

 

[Studiot]

I have now had a chance to read the refence from Wooton et al linked to by Sylas.

Many sources now state that ocean pH has already changed 0.1 units over the past century (3, 10). The basis for these statements is model simulations that include only physical processes in the control of pH and that are calibrated from a single year of data (3) rather than those that use direct empirical measurements of ocean pH through time. Little published empirical information exists on the dynamics of directly measured ocean pH (1, 11), and none is available at temperate latitudes, which harbor the world's most productive fisheries.

So the authors admit that there is a paucity of data and commendably set about gathering some.

Their results certainly demonstrate changeability and a significant difference from the above extract from their opening statement compared with conventional wisdom about buffering in the pH range 7 - 9.

Although their results are taken over 8 years, rather better than the other results they refer to, this is a very short time to distinguish between a genuine downward trend and the catching the downward part of a cyclic phenomenon.

Further they attribute the lowering of pH to increasing atmospheric concentration of CO₂ but I could not find any matching measurements to substantiate or deny this theory.

Conventional chemical wisdom is that there is buffering action on pH, via the bicarbonate ion, when you add CO₂ to seawater, even in substatial quantities. This is because the bicarbonate ion can participate in two different reactions one producing pH reducing protons and the other producing pH raising hydroxyl ions.
So the conventional expectation to change the pH is by the introduction of another protonic acid.

I will post all the reactions if anyone is sufficiently interested, they are quite a complex set.

The authors confirm this chemistry later in the article but do not appear to have studied whether any of the species ( or other input) may have produced such an acid.

I agree the figures are a bit of a puzzle and beg considerable further investigation, especially to see if they can be reproduced elsewhere in the world.

In answer to the original question

Conventional chemistry suggests that CO₂ will only be released if the pH falls below about 4. Interestingly this is the pH at which concrete loses its ability to protect reinforcement from corrosion.

 

[sylas]

pH measures both acidity and alkalinity. A low Ph number indicates acidity and a high pH rating indicates alkalinity.

To get more specific, there are differnt ways to measure ocean water.

No; a high pH indicates basicity, which is not the same as alkalinity. They are different things to measure, not two ways to measure the same thing.

http://en.wikipedia.org/wiki/PH#Mathematical_Definition

If we are using wikipedia, then look up the definition of alkalinity, not the definition of pH! You are looking up the wrong thing to understand what the paper is actually talking about. From Alkalinity:

Alkalinity is sometimes incorrectly used interchangeably with basicity. For example, the pH of a solution can be lowered by the addition of CO2. This will reduce the basicity; however, the alkalinity will remain unchanged (see example below).​

So there is precedent for your usage, but it is not actually technically correct and in particular you should use the more correct definitions when reading the paper cited.

A more comprehensive reference (which is given in the wikipedia article) is http://132.239.122.17/co2qc/handbook.html, which gives more formal definitions. Alkalinity in seawater is defined there as a sum of concentrations of proton acceptors over proton donors and free hydrogen ions. There is an equation in SOP 3 at the above link which is basically listing the relevant ions.

This research really isn't conflicting at all with the information I have given previously of more direct relevance to the original question; although it certainly is an interesting addition to understanding of the inorganic carbon cycle.

Cheers -- sylas

PS. Missed a flurry of posts while writing. I think we are almost on the same page now with respect to pH and alkalinity -- except that pOH is not alkalinity either... it is still basicity. (I'll defer to Borek with any chemistry questions, by the way!) But the important point is that the research cited on fish and carbonates isn't in fact in any conflict with answers given for the original question, about the relationship between pH and CO2 being absorbed or emitted from seawater.

 

[Evo]

PS. Missed a flurry of posts while writing. I think we are almost on the same page now with respect to pH and alkalinity -- except that pOH is not alkalinity either... it is still basicity. (I'll defer to Borek with any chemistry questions, by the way!) But the important point is that the research cited on fish and carbonates isn't in fact in any conflict with answers given for the original question, about the relationship between pH and CO2 being absorbed or emitted from seawater.

Yes, I was editing, sorry. You are correct about the relationship between CO2 and acidity, but if the study in Science is correct, the fish are expected to increase alkalinity.

 

[sylas]

That's a good point, but I don't claim to know if that is true for sure.

What I do know is that as water warms, its ability to hold dissolved gases, such as CO2 is diminished. So, theoretically, atmospheric CO2 levels do not need to fall for the oceans to release CO2. If the oceans were to warm enough, they could begin to release CO2 even if atmospheric CO2 levels were to continue to rise. It's all a matter of the relative response of each characteristic.

The essential sentence here is the last. The chemical reactions will drive things towards the natural equilibrium levels. With the current huge excess of atmospheric carbon, that means the ocean will be a net absorber.

If there was no geological carbon being added to the atmosphere and the whole system settled to an equilibrium, then atmospheric levels of carbon would fall, initially at roughly the rate it is currently rising, as carbon continues to be taken up into the ocean and land sinks, but the effect of a warmer ocean would be that the final equilibrium would be with atmospheric levels a bit higher that the natural pre-industrial background of 280 ppm, but a lot lower than the current 390 or so. There's way too much atmospheric carbon at present for the reactions to be pushing the ocean to a net emitter.

Cheers -- sylas

 

[sylas]

Yes, I was editing, sorry. You are correct about the relationship between CO2 and acidity, but if the study in Science is correct, the fish are expected to increase alkalinity.

Sure; I'm with you there.

 

[Borek]

Conventional chemical wisdom is that there is buffering action on pH, via the bicarbonate ion, when you add CO₂ to seawater, even in substatial quantities. This is because the bicarbonate ion can participate in two different reactions one producing pH reducing protons and the other producing pH raising hydroxyl ions.
So the conventional expectation to change the pH is by the introduction of another protonic acid.

I have a feeling that you have ignored fact that to get bicarbonate from carbon dioxide you have to actually produce H⁺, lowering pH before the buffering effect appears.

Conventional chemistry suggests that CO₂ will only be released if the pH falls below about 4.

Can you explain what is justification (is that a correct word?) of that number?

What happens to carbon dioxide (whether it leaves solution or dissolves) depends not only of solution pH, but also on partial pressure of CO₂ above the solution. In a theoretical case where you start with sodium carbonate and you put the solution in an carbon dioxide free atmosphere, after time long enough you are left with solution of just NaOH, all carbon dioxide leaves the solution. In practice it doesn't happen as partial pressure of carbon dioxide present in the atmosphere is high enough so that CO₂ actually dissolves. But from what I remember pH of pure water saturated with atmospheric carbon dioxide falls to about 5.5 or 5.6, any solution with pH below that limit should be releasing CO₂.

 

[sylas]

But from what I remember pH of pure water saturated with atmospheric carbon dioxide falls to about 5.5 or 5.6, any solution with pH below that limit should be releasing CO₂.

And, just for the record, the ocean is not likely to get down to pH values like that. Yikes.

Ocean "acidification" is a potentially misleading term. The pH of sea water is above 7; sea water is slightly alkaline. The phrase "acidification" simply means that measured pH is, on average, falling.

Also: precise equilibrium chemistry levels are not entirely worked out, but it is widely believed that the large rises in CO2 from the depths of the ice age up to the balmy conditions of the Holocene (180 ppm to 280 ppm) is because CO2 was released from the ocean as it warmed.

 

[Borek]

Ocean "acidification" is a potentially misleading term.

Not for me

 

[Studiot]

C{O_{2(gas)}} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} C{O_{2(aq)}}

{H_2}O + C{O_{2(aq)}} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H_2}C{O_3}

{H_2}C{O_3} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H^ + } + HCO_3^ -

HCO_3^ - \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H^ + } + CO_3^{2 - }

CO_3^{2 - } + {H_2}O \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} HCO_3^ - + O{H^ - }

HCO_3^ - \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} C{O_{2(aq)}} + O{H^ - }

 

[Borek]

Actually if you add

H₂O <-> H⁺ + OH^-

last two equations you have listed become redundant. System is fully described by 5 equilibrium constants.

 

[sylas]

Actually if you add

H₂O <-> H⁺ + OH^-

last two equations you have listed become redundant. System is fully described by 5 equilibrium constants.

... and, for masochists or Chemistry Gods, you can look up typical values for those constants in the handbook reference I linked in [post=2690047]msg #25[/post]: http://132.239.122.17/co2qc/handbook.html, with suggestions for solving the equations. Values will vary with conditions in different parts of the ocean, of course, but over all: more CO2 absorbed leads to lower pH.

 

[Studiot]

Actually if you add

H2O <-> H+ + OH-

Not really.

Without the other ions being present the dissociation constants are quite different.

The situation is further complicated by all the other ions p[resent in seawater.

It was even more complicated by trying to paste from Mathtype, just now.

Also these equations carry the reason for the buffering at pH = 8.8 and again at pH = 4.5, which was the other question asked about my posts.

 

[Borek]

Not really.

Without the other ions being present the dissociation constants are quite different.

Can you give example of what you mean? Assuming - for now - that there is no other acid/base systems present water/carbon dioxide system contains H₂O, H⁺ (well, no such thing as a free proton, but let's not complicate things further), OH^-, CO₂, H₂CO₃, HCO₃^- and CO₃^2-. You have listed 6 reactions between them - it is possible to list further ones, but only 5 are independent. For example you have listed:

HCO₃^- <-> H⁺ + CO₃^2-

and

CO₃^2- + H₂O <-> HCO₃^- + OH^-

When you take water autodissociation into account, this is in fact the same equilibrium (that's what Brønsted-Lowry theory is about). Acid dissociation constant for the first reaction is

K_{a2} = \frac {[H^+][CO_3^{2-}]} {[HCO_3^-]}

base dissociation constant for the second is

K_{b1} = \frac {[HCO_3^-][OH^-]} {[CO_3^{2-}]}

(water concentration is assumed to be constant and ignored by convention)

and

K_{a2} K_{b1} = K_w

where K_w is water ionization constant.

The situation is further complicated by all the other ions p[resent in seawater.

Yes, but we can ignore them for now to make sure we talk about the same system.

Also these equations carry the reason for the buffering at pH = 8.8 and again at pH = 4.5, which was the other question asked about my posts.

I can't see it, please elaborate. Carbonates have a significant buffering effect at pH around 6.37 (pK_a1) and 10.25 (pK_a2). See attached image. This is a plot of a carbonate buffer capacity, horizontal scale is pH, black arrow points at pH 8.8 (where buffering capacity is relatively low). For low pH this plot assumes high pressure of carbon dioxide, but it is irrelevant in the range we are talking about.

--

 

[Studiot]

I will post some more theory as soon as I get a chance,
meanwhile I offer this experiment.

Take a beaker of water, with a pH meter inserted. What is the pH?

Stir in a spoonful of common salt what now is the pH?

Stir in a spoonful of bicarbonate of soda what now is the pH?

Add a knob of dry ice what now is the pH?

 

[Borek]

Results will depend on water source. Assuming for simplicity we start with a distilled water - at first it will be slightly acidic, something around 6 or even less, as atmospheric carbon dioxide dissolves pretty fast and keeping water pure is very difficult. Addition of NaCl - increase of ionic strength will shift pH up a little bit, but change will be below 0.1 pH unit. Bicarbonate - pH goes up, how far depends on concentration, we are talking about pH 8 range. Dry ice - pH goes down, but it is rather impossible to calculate how far, as this is a dynamic situation and final concentration of dissolved gas depends on many factors.

 

[aquitaine]

How can the ocean simultaneously release CO2 and decrease in pH?

Currently the ocean is acidifying, as it absorbs about a third of the fossil-carbon dioxide that we emit, which then in part assumes the form of carbonic acid. But in the future, if the increasing atmospheric greenhouse effect continues to also warm the ocean enough, we expect this absorption will be reversed and vast quantities of CO2 will distil out from the ocean.

Nonetheless, apparently we do not expect warming to cause any reverse to the acidification. (I asked one paleoceanographer/marine-chemist, and heard there is no contradiction for water to be simultaneously decreasing in pH and liberating CO2.) But naively, if warmed water begins losing carbon, then shouldn't the concentration of carbonic acid fall (and hence the pH start to rise back again)?

Edit: The topic of this thread is not "global warming or climate change". It is purely an ocean-chemistry question. Regardless of what is actually happening to our ocean (or rather, regardless of what external factors may be controlling the temperature of and the partial pressure of CO2 above a hypothetical test-ocean) the question is simply whether in principle such an ocean hypothetically could ever be driven (by adjusting those two parameters) to release CO2 while simultaneously to decrease in pH, and how exactly? (So this is what self-censorship is like..)

A little late to this topic, but perhaps something to consider looking into is how does the current ocean pH now compare with 50+ million years ago, back when CO2 levels were significantly higher than today (and there were no polar ice caps at all)? Perhaps it could be a good insight as to where the acidification issue is going as CO2 levels rise. Just a thought.

 

[sylas]

A little late to this topic, but perhaps something to consider looking into is how does the current ocean pH now compare with 50+ million years ago, back when CO2 levels were significantly higher than today (and there were no polar ice caps at all)? Perhaps it could be a good insight as to where the acidification issue is going as CO2 levels rise. Just a thought.

I'm not familiar with that work. There are attempts to estimate pH in the past; but there are all kinds of additional uncertainties. I had a quick look and do not feel competent to summarize all the questions and variables for such reconstructions.

You are quite right that paleo studies are an important part of the whole picture and a useful test for various theories, but in general, my impression is for the most part, that we get a better view of what is going on in the present and in the immediate future from more direct data, and then this helps us interpret the records from the past.

Cheers -- sylas

 

[Studiot]

drat, I was going to add numbering to my list of equations so I could refer to them, but can't seem to edit that post now.

 

[Evo]

drat, I was going to add numbering to my list of equations so I could refer to them, but can't seem to edit that post now.

Just paste the corrected version into another post and I can edit your old post, if you'd like.

 

[Studiot]

Yes please any numbering you can do would be great.

 

[Evo]

Yes please any numbering you can do would be great.

Just copy your old post into a new post, edit it, then I will move it back into the old post for you.

 

[aquitaine]

I'm not familiar with that work. There are attempts to estimate pH in the past; but there are all kinds of additional uncertainties. I had a quick look and do not feel competent to summarize all the questions and variables for such reconstructions.

You are quite right that paleo studies are an important part of the whole picture and a useful test for various theories, but in general, my impression is for the most part, that we get a better view of what is going on in the present and in the immediate future from more direct data, and then this helps us interpret the records from the past.

Cheers -- sylas

Just a coincidence but this just happened to come up in physorg recently.

In a paper published April 26 in the Proceedings of the National Academy of Sciences, a team of researchers led by a Stanford geologist said that as carbon dioxide gas dissolved in the oceans, it raised the acidity of seawater.

The research team said it was a deadly combination - carbon dioxide in the atmosphere and higher acidity in the oceans - that eventually wiped out 90 percent of marine species and about three-quarters of land species, in a cataclysmic event 250 million years ago known as the "end-Permian extinction."

Back then, the ocean teemed with corals, algae, clams and snails. Soon after, however, there was an abrupt change to a thick layer of bacteria and limestone, a "slime-world," dominated by bacteria.

Just in case you change your mind. :)

 

[Gokul43201]

Here:

C{O_{2(gas)}} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} C{O_{2(aq)}} ~~~-[1]

{H_2}O + C{O_{2(aq)}} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H_2}C{O_3}~~~-[2]

{H_2}C{O_3} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H^ + } + HCO_3^ -~~~-[3]

HCO_3^ - \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {H^ + } + CO_3^{2 - }~~~-[4]

CO_3^{2 - } + {H_2}O \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} HCO_3^ - + O{H^ - }~~~-[5]

HCO_3^ - \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over<br /> {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} C{O_{2(aq)}} + O{H^ - }~~~-[6]

No need to edit old posts.

Also, call dissociation of water into (for simplicity) H+ and OH- equation [7].

Equations 4,5 and 7 involve a redundancy (4+5=7), as do equations 2,3,6 and 7 (2+3+6=7). Therefore you are left with no more than 5 independent equations.

 

[Studiot]

Many thanks Gokul for doing that, as well as to Evo for offering. I did try last night but got really tangled up.

Anyway the reason for posting the list of equations is that we are trying to discuss the acidification of at least the surface layers of the ocean.

Also there is another measured parameter, the alkalinity, which is a measure of the ability of the water sample to act a base by reacting with protons.

An apparent paradox occurs in this system as both the acidity (pH) and the alkalinity can simultaneously increase.

To discuss all this we need some to understand how the acidity (pH) and alkalinity is measured, along with how that affects the equilibrium position of the reactions. The point being that natural waters can act as either a base or an acid according to circumstance.

It is not possible to determine the concentrations of the reactants without disturbing the reactions so the standard methods involve titrating to an endpoint.
The standard endpoint indicators in use are either Methyl Orange or Phenolphthalein.
The former changes colour at pH = 4 and the latter at pH = 10 so the choice of titration indicator will influence the results.

A plot of a typical titration is attached.

Additionally whilst the bicarbonate reactions are very fast in the lab, ocean mixing times mean that the response of a large inhomogeneous mass of water is much slower.

With reference to my equations it is important not to combine them as has been suggested, none are ‘redundant’.

Equations 3 and 4 form the reactions of the so called ‘carbonate buffer’.

Equation 4 is the first of two steps in the formation of carbonic acid: the carbonate ion in solution removes one proton forming the bicarbonate ion. Thus the pH rises or is prevented from falling if protons are being added.

Equation 3 is the second step in the formation of carbonic acid by removal of another proton.

Now these reactions occur around points C and D of the titration curve. Thus a titration with methyl orange as an indicator will reach DE on the curve at a pH of 4-5.
This corresponds to an endpoint where most of the bicarbonate and carbonate has been converted to carbonic acid.
Whilst a titration with phenolphthalein as indicator will end somewhere on BC and corresponds to and endpoint for the situation where most of the carbonate ion has been converted to bicarbonate but little of the bicarbonate has been converted to carbonic acid.


This is the basic chemistry of the ‘buffer’.
Most seawater is actually in contact with calcium carbonate and other ion sources, which replenish the participants in the carbonate buffer.
Set against this is the entry of carbon dioxide from the atmosphere.

I will discuss how that affects the situation in the light of the other equations, in my next post.

 

[Borek]

With reference to my equations it is important not to combine them as has been suggested, none are ‘redundant’.

No. Equilibrium in this system (water/carbon dioxide) is described by 5 equilibrium constants. You may write as many reactions as you want, but equilibrium constant for every other reaction but the five basic ones can be calculated from these initial ones. In this sense every other reaction is redundant.

I will give another example. Let's concentrate on just acid dissociation for a moment. Carbonic acid dissociates in two steps:

H₂CO₃ <-> H⁺ + HCO₃^-

and

HCO₃^- <-> H⁺ + CO₃^2-

with respective dissociation constants

K_{a1} = \frac {[H^+][HCO_3^-]} {[H_2CO_3]}

and

K_{a2} = \frac {[H^+][CO_3^{2-}]} {[HCO_3^-]}

Someone may say these are not all reactions describing dissociation of carbonic acid, as there is also overall dissociation reaction

H₂CO₃ <-> 2H⁺ + CO₃^2-

with overall dissociation constant

K_{a12} = \frac {[H^+]^2[CO_3^{2-}]} {[H_2CO_3]}

However, this part of the system is NOT described by three reactions and three constants, as

K_{a12} = \frac {[H^+]^2[CO_3^{2-}]} {[H_2CO_3]} = \frac {[H^+][HCO_3^-]} {[H_2CO_3]} \frac {[H^+][CO_3^{2-}]} {[HCO_3^-]} = K_{a1}K_{a2}

- so this reaction and its equilibrium constant doesn't add new information about the system. Knowing any two of three these constant - K_a1, K_a2, K_a12 - we can calculate third one. That means one of these reactions is redundant. Same logic can be applied to the reactions you have listed. There are five independent equilibrium constants describing full system, if you have more than five reactions - some of them are redundant, as their equilibrium constants can be calculated from the basic 5.

Equations 3 and 4 form the reactions of the so called ‘carbonate buffer’.

I would say these are two independent buffers, based on two different acid/conjugate base pairs.

Equation 4 is the first of two steps in the formation of carbonic acid: the carbonate ion in solution removes one proton forming the bicarbonate ion. Thus the pH rises or is prevented from falling if protons are being added.

Equation 3 is the second step in the formation of carbonic acid by removal of another proton.

Now these reactions occur around points C and D of the titration curve.

No. Reaction 3 takes place between points A and B, while reaction 4 takes place between points C and D.

I happen to have a little bit better version of the titration curve you refer to:

(used by permission, see www.titrations.info/acid-base-titration-polyprotics-and-mixtures)

This is not exactly the same situation, as this is titration curve for Warder titration - so there is not only carbonate, but also NaOH present initially, hence end point at pH 11.34 - but after that moment it is the titration you are talking about, even with color ranges marked.

 

[Studiot]

No. Equilibrium in this system (water/carbon dioxide) is described by 5 equilibrium constants

Actually I wasn't describing the carbon dioxide/water system.
I was describing the actual chemical mix you are likely to find in natural waters, before introducing carbon dioxide.

I probably didn't present the equations in the most sensible order and I apologise for this.

The carbonate ions are present from other sources and the question is what happens if we now introduce carbon dioxide?
Also what happens if there is replenishment of carbonate?

I would say these are two independent buffers, based on two different acid/conjugate base pairs.

Yes that is true but I am only using the name granted to the pair of buffer reactions by all the authorities I have read.

No. Reaction 3 takes place between points A and B,

The natural waters I have described cannot attain the high pH between A an B so neither reaction can be here.

I think you have reactions 3 and 4 the wrong way round, 4 must come before 3 as you add protons.

I agree with your concatenation of reaction constants, but have not yet reached the point where they are relevant.

 

[Borek]

Actually I wasn't describing the carbon dioxide/water system.
I was describing the actual chemical mix you are likely to find in natural waters, before introducing carbon dioxide.

Yes and no. "Natural mix" would contain many different ions as well (with their own buffering capabilities, think borates, phosphates, ammonia and so on, even humic acids in some cases), so far you have concentrated on system built around carbonates equilibria - and technically any system containing carbonates already contains carbon dioxide. When you add CO₂ you will be just shifting equilibria present, but the chemistry will be the same.

I probably didn't present the equations in the most sensible order and I apologise for this.

Order is not a problem (well, obviously it is - I have mistaken numbers when referring to the reactions and titration curve). Selection of the reactions - is. And you should really add water autodissociation.

The carbonate ions are present from other sources and the question is what happens if we now introduce carbon dioxide?
Also what happens if there is replenishment of carbonate?

What is the source of carbonates? If it is calcium/magnesium carbonate, in both cases pH goes down, although presence of carbonates slows the process down. And if that's the case pH goes down and alkalinity goes up, I have signalled before that there is no simple dependency between both.

Yes that is true but I am only using the name granted to the pair of buffer reactions by all the authorities I have read.

As far as I can tell one of these systems is called carbonate buffer, the other bicarbonate buffer. Probably calling them collectively "carbonate buffers" is OK, my English ducked behind the table and pretends to be not here.

The natural waters I have described cannot attain the high pH between A an B so neither reaction can be here.

OK, but when you state reaction 4 takes place at C-D it is confusing. After first end point concentration of CO₃^2- is neglectable. At the first end point CO₃^2- constitutes about 1% of all forms of carbonates present. This part of the titration curve (C-D) is dominated by the protonation of bicarbonate (or by the bicarbonate buffering).

I think you have reactions 3 and 4 the wrong way round, 4 must come before 3 as you add protons.

Yes, sorry - my mistake.

--