What Keeps Sea Water Alkaline Despite the Presence of Carbon Dioxide?

[vivesdn]

Several marine organisms including coral and shelfish create calcium and magnesium carbonate structures.
As carbonate ion comes from dissolved CO2, this process should acidify the ocean. But after millions of years, the ocean still is slightly alkaline.
What is the process that consumes acid excess or that generates alkali to keep sea water alkaline?

 

[Studiot]

For what it's worth this thread contains the chemistry of the acidity / alkalinity of seawater.

https://www.physicsforums.com/showthread.php?t=398208&highlight=seawater

go well

 

[olivermsun]

It may be helpful to understand that calcium carbonate is not the same as dissolved CO2. In fact, most "CO₂" in the ocean exists neither as dissolved CO₂(aq) nor as carbonate, CO₃^2-. Instead, most exists as the bicarbonate ion HCO₃^-.

The equilibrium reactions between these 3 constituents determine their relative abundance in seawater. There is always plenty of bicarbonate (and lots of other stuff as well) washing down into the oceans, e.g., from weathering of rocks on land. In addition, dissolved CO₂(aq) is roughly in equilibrium with the atmosphere, since gas exchange can take place through the air-sea surface.

The previous thread linked by Studiot does a much more detailed job of explaining the chemistry. It is also pointed out in the thread that seawater alkalinity is not the "opposite" of pH. You could very well have the same alkalinity even while pH declines due to increasing atmospheric pCO₂.

 

[morrobay]

Well like one of my professors at San Diego State said ;

"You can make carbonate chemistry do just about anything you want it to do "

 

[vivesdn]

In fact, most "CO₂" in the ocean exists neither as dissolved CO₂(aq) nor as carbonate, CO₃^2-. Instead, most exists as the bicarbonate ion HCO₃^-.

But organisms create structures of calcium and magnesium carbonate. As the three forms are in equilibrium this process wil acidify sea water.

And I'm just asking what is the counterpart of this process as sea water is not acidic.

If the answer is the weathering of rocks, then you should explain how those rocks were formed without lowering the pH of the ocean.

 

[olivermsun]

Why would fixing of carbonate into calcium carbonate acidify water when it removes dissolved inorganic carbon from the water column?

 

[Studiot]

As the three forms are in equilibrium this process wil acidify sea water.

What makes you think they are in equilibrium?

Surely the point is that they are not in equilibrium so driving the coupled reactions in one direction or the other making the resultant pH whatever it end up.

 

[SpectraCat]

Why would fixing of carbonate into calcium carbonate acidify water when it removes dissolved inorganic carbon from the water column?

Because when you form carbonate from the reaction of CO2 with water, you end up releasing two protons. Thus the OP is technically correct ... however, the extent to which this happens in the ocean at pH~7.9 is very small. Still, the OP is also correct that if the organisms did just scavenge already formed carbonate ions from ocean water, it would stress the coupled equilibria such that more CO2 would be absorbed, resulting in further acidification (assuming that these were the only chemical equilibria in the system).

My guess (and it is only a guess) is that the organisms actually use enzymes to generate carbonate from bicarbonate metabolically, and then the extra protons are dealt using some sort of biological buffer system within the organism, and not released back into the ocean.

Another possibility is that there is a buffer system (or more than one) so that the pH changes are balanced, however that just shifts the OP's questions to "what are those buffer systems and how have they remained in balance over milliions of years". Perhaps that is a more appropriate question, but I am not sure.

 

[olivermsun]

You and the OP are both right -- marine organisms fix bicarbonate into calcium carbonate shellsin the euphotic zone, releasing H+ and CO2 (which should in turn slow down dissolution of CO2 in the upper ocean?). However the inorganic equilibrium carries CaCO3 downward, where it eventually undergoes the reverse reaction (dissolution, same reaction as in limestone weathering) and releases CO2 at depth. Hence a CO2 "buffer" by transport into the deep ocean.

Another possibility is that there is a buffer system (or more than one) so that the pH changes are balanced, however that just shifts the OP's questions to "what are those buffer systems and how have they remained in balance over milliions of years". Perhaps that is a more appropriate question, but I am not sure.

I am not sure either. If the question remains why pH has remained relatively steady over long time periods, then it could be said that seawater alkalinity IS the buffer, and also that atmospheric pCO2 has also been (relatively) stable in recent epochs. Perhaps the OP would further clarify?

 

[vivesdn]

I've visited several mountain ranges that are formed of limestone. Millions (billions, trillions) of tons of carbonate out of the water.

It must exist a process to counterpart the protons released by the organisms. It cannot be a buffer as with time it will also shift. Except if that buffer is open (for example, released on vents).
Maybe the organisms themseleves are consuming the protons as pointed by SpectraCat?
From what I've read, photosynthesis does not consume the protons but I may be wrong...

 

[Borek]

I wonder if that's really a problem. Both Ca and C circulate in the environment, being deposited as CaCO₃ (and in many other forms), then freed out (by weathering, volcanic processes and so on). Could be pH that we observe is already a steady state (in other words - oceans are already acidified to the equilibrium point).

 

[morrobay]

The other side of this question is what happened to the anion associated with the
Ca⁺⁺ Suppose that anion was hydroxide, Ca⁺⁺ (OH^-)₂ lime water. So:
2 Ca⁺⁺(OH^-)₂ + 2H₂CO₃ >>> shellfish >> 2CaCO₃ + 4 H₂O or
Ca(OH)₂ + CO₂ >>shellfish> CaCO₃ + H₂O

 

[Borek]

Suppose that anion was hydroxide

Suppose it was not? Composition of anions in sea water is quite stable and known.

 

[morrobay]

The other side of this question is what happened to the anion associated with the
Ca⁺⁺ Suppose that anion was hydroxide, Ca⁺⁺ (OH^-)₂ lime water. So:
2 Ca⁺⁺(OH^-)₂ + 2H₂CO₃ >>> shellfish >> 2CaCO₃ + 4 H₂O or
Ca(OH)₂ + CO₂ >>shellfish> CaCO₃ + H₂O

Hydroxide is one of the Ca⁺⁺ anions among phosphate, sulfate, chloride...
Calcium dissolved from rocks, where some shellfish attach, reacts with water:
Ca + 2 H₂O ---> Ca(OH)₂ + H₂
In the previous post, the reaction with carbonic acid and calcium hydroxide is a possible
reaction that produces calcium carbonate for the shellfish as well as consuming H⁺

 

[vivesdn]

produces calcium carbonate for the shellfish as well as consuming H+

Wrong: it consumes hidroxide anion if this is the anion, or produces protons if the anion was cloride, phosphate, sulphate or whatever.

Are we observing the steady state pointed by Borek? Was the ancient sea far more alkaline? Or is there a process that consumes the acid?

 

[Borek]

Calcium dissolved from rocks, where some shellfish attach, reacts with water:
Ca + 2 H₂O ---> Ca(OH)₂ + H₂

This is completely off and suggests you have no idea what you are talking about. Calcium in rocks is always in the ionic form, metallic calcium is very reactive and never present in native form in nature.

 

[morrobay]

Several marine organisms including coral and shelfish create calcium and magnesium carbonate structures.
As carbonate ion comes from dissolved CO2, this process should acidify the ocean. But after millions of years, the ocean still is slightly alkaline.
What is the process that consumes acid excess or that generates alkali to keep sea water alkaline?

Have you forgotten what your original question was ?
That I just answered for you.
Not talking about chlorides , not talking about steady states. Just the following reaction
that produces calcium carbonate and consumes the hydrogen ions from the carbonic acid:
2Ca(OH)₂ + 2H₂CO₃ ---> 2CaCO₃ +4H₂O.
Do you understand that on the right side that calcium carbonate is a product and the
the 4 waters are a product , from 4H⁺ and 4OH^-
That the 4H⁺ , from the carbonic acid are consumed ?
If not you should take a class in introductory chemistry

 

[Borek]

2Ca(OH)₂ + 2H₂CO₃ ---> 2CaCO₃ +4H₂O

You are assuming here calcium in water is present in the form of hydroxide, but that's not true.

Sea water contains a lot of cations and anions, between them Ca²⁺, HCO₃^- and CO₃^2-. At pH around 8 (which is more or less sea water pH) HCO₃^- is a dominating form (well over 90% of dissolved carbonates). That means when calcium carbonate is removed from the sea water you are removing Ca²⁺ and HCO₃^-, and one of the products is H⁺:

Ca²⁺ + HCO₃^- -> CaCO₃ + H⁺

That's the reaction OP refers to.

 

[morrobay]

This is completely off and suggests you have no idea what you are talking about. Calcium in rocks is always in the ionic form, metallic calcium is very reactive and never present in native form in nature.

Yes I am assuming that calcium hydroxide exists in seawater, for example:
CaO + H₂O --> Ca(OH)₂
And it is this calcium hydroxide that I have shown several times as a reactant in the reaction that produces calcium carbonate and maintains alkalinity.
And in the history of calcium hydroxide in seawater are you sure that none of it was produced
from Ca + H₂O ?
Are there any Geo - Marine Chemists out there to give a definitive answer here ?

 

[Borek]

Yes I am assuming that calcium hydroxide exists in seawater

Think about it this way - at pH 8 [OH^-] = 10^-6M, while concentration of calcium in sea water is around 410 ppm, that is about 0.01M - 10⁴ times larger than [OH^-]. Doesn't make sense to speak about Ca(OH)₂.

And in the history of calcium hydroxide in seawater are you sure that none of it was produced from Ca + H₂O ?

Yes I am sure, for the reasons stated earlier.

 

[morrobay]

Think about it this way - at pH 8 [OH^-] = 10^-6M, while concentration of calcium in sea water is around 410 ppm, that is about 0.01M - 10⁴ times larger than [OH^-]. Doesn't make sense to speak about Ca(OH)₂.

This is correct for in solution calcium hydroxide . But what is not being counted here is the
solid, precipitated form of calcium hydroxide from lime water ,
CaO + H₂O ---> Ca(OH)₂
If this form of calcium hydroxide is available to shellfish then it does make sense.

 

[Borek]

If this form of calcium hydroxide is available to shellfish then it does make sense.

But it isn't. All they have is a sea water and ions dissolved.

No matter how you will try you will not add sense to your posts about Ca(OH)₂. It is a waste of time.

 

[morrobay]

But it isn't. All they have is a sea water and ions dissolved.

Yes and the solubility of those dissolved ions and the precipitate are in dynamic
equilibrium: Ca(OH)₂ >> << Ca⁺⁺ + 2OH^-
as a function of pH, temperature and common ions.
Limestone is the major submarine geologic structure in parts of the ocean. Shellfish
inhabit the marine sediments in those areas.

 

[morrobay]

But it isn't. All they have is a sea water and ions dissolved.

Yes and the solubility of those dissolved ions and the precipitate are in dynamic
equilibrium: Ca(OH)₂ >> << Ca⁺⁺ + 2OH^-
as a function of pH, temperature and common ions.
Limestone is the major submarine geologic structure in parts of the ocean. Shellfish
inhabit the marine sediments in those areas.

Add: You can apply the Le Chatelier principle here, as the shellfish consume Ca and OH ions in
marine sediments that contain precipitated Ca(OH)₂ the above reaction
goes to the right. So the shellfish have available sources of Ca and OH ions that are not
reflected in Ca ppm or overall pH.
Once again I am not stating that this reaction: 2 Ca(OH)2 + 2H2CO3 --> 2 CaCO3 + 4H2O
is the only way shellfish make calcium carbonate.

 

[Borek]

Be warned - now you are just trolling. There is no Ca(OH)₂ in the sea precipitates, limestone doesn't contain it either, shellfish doesn't use solid as a source of materials needed for CaCO₃. As I told you - no matter how you will try, you are not going to make situation look better for you, you don't understand chemistry behind.

 

[vivesdn]

2Ca(OH)2 + 2H2CO3 ---> 2CaCO3 +4H2O.
Do you understand that on the right side that calcium carbonate is a product and the
the 4 waters are a product , from 4H+ and 4OH-
That the 4H+ , from the carbonic acid are consumed ?

And OH- is also consumed, didn't you notice?

 

[morrobay]

Here is a reaction that removes calcium carbonate from seawater without
producing H⁺

Ca⁺⁺ + 2HCO_3^- --> Ca(HCO₃)₂
--> CaCO₃ +H₂O + CO₂

 

[Borek]

Here is a reaction that removes calcium carbonate from seawater without
producing H⁺

Ca⁺⁺ + 2HCO_3^- --> Ca(HCO₃)₂
--> CaCO₃ +H₂O + CO₂

No. You liberated carbon dioxide, it reacts with water acidifying solution. Initially solution pH was around 8.3, now it is somewhere between 5 and 6.

 

[morrobay]

Yes it can do both: cold temperatures favor CO₂ + H₂O -->
2H⁺ + CO_3^-- So your right with some local conditions.

But with warmer temperatures the CO₂ (gas) can be released to atmosphere

 

[Borek]

Yes it can do both: cold temperatures favor CO₂ + H₂O -->
2H⁺ + CO_3^-- So your right with some local conditions.

But with warmer temperatures the CO₂ (gas) can be released to atmosphere

Looks unrealistic. Solubility of carbon dioxide at 15 °C is 0.197 g, at 35 °C is 0.111 g (per 100g of water, 760 mmHg of gas pressure), that means it changes about twice in the range of reasonable water temperatures. Twofold change of acid concentration means pH change in the range of several tenths of unit (below 0.5) - we are still on the acidic side.