Acquired Reading

Shorter than a tweet, but an excellent review of ocean acidification from the National Academy.

Bio types can follow on to the next chapter.

Enjoy

Of course, this arose in one of those Dunning-Krugar things about acidification of the oceans, like how can you talk about acidification when then pH of the oceans is ~8.  Eli has had some words about that in the past (see also the comments by Kenneth Johnson and Bernard J).

The root of teh probelm dates back to Arrhenius (yes, him again) who was the first to define the chemical basis of acids and bases.  Since the Earth is a water world and in 1884 organic chemistry was the wild west and the properties of ions in solutions newfangled stuff that would win him a Nobel Prize in 1903, Svante A defined an acid as a source of H⁺ ions and a base as a source of OH^- ions.  The problem is that there are other ions which are the sources of both via the hydrolysis of water.  For example, if you dump a bunch of carbonate (CO3)^2- ions into water you get OH^- via

H2O(l) + CO3(aq)^2-  -->  HCO3(aq) ^- + OH(aq)^-

Using the Arrhenius definition of acids and bases quickly gets a bunny into trouble when discussing what happens in complex aqueous solutions where there are a variety of ions, let alone solutions in solvents besides water, a trouble that was ameliorated by Brønsted and Lowry, who moved the attention of chemists to the behavior of hydrogen ions aka protons. 

In their view an acid was any molecule or ion that donates a hydrogen ion to another and a base any species that receives it.  In the reaction above, chemists would describe the water molecule (H2O(l)) as an acid because one of its hydrogens is donated to the carbonate ion CO3^2-to form the hydrogen carbonate ion  HCO3 ^-on the product side.  Similarly, the carbonate ion CO3^2-  is a base, because it receives the hydrogen ion.   

In the reverse reaction the OH(aq)^- accepts a proton so it is the base and the HCO3(aq) ^- donates one so it is an acid. 

From the Brønsted Lowry point of view, OH(aq)^- is nothing special, just another damned proton catcher. 

ADDED:  A useful example of this is neutralization of carbonic acid by carbonate ions.  Carbonic acid is formed when CO2(aq) reacts with water

H2O(l) + CO2(aq) -->  H2CO3(aq)

 

 The carbonic acid then can react with the carbonate ion to form two hydrogen carbonate ions

 H2CO3(aq) + CO3(aq)^2-  -->  HCO3(aq)^-  + HCO3(aq) ^-

there is no OH(aq)^- generated in this reaction but the carbonic acid is neutralized by the  CO3(aq)^2- which is the base.  In the reverse reaction one of the hydrogen carbonates is an acid (proton doner) and the other a base (proton acceptor).  FWIW water and  HCO3(aq) ^- can both catch and toss protons, so they are called amphoteric.
 
From this point of view alkalinity is defined as the capacity to neutralize acid, or if you will to catch protons.  

On the water world, this makes sense because the lakes, streams and oceans are filled with ions of weak acids like carbonic acid, which can hydrolyze water.  That is why alkalinity and especially the alkalinity of the oceans is defined as the concentrations (indicated by [] and see the link to the NAS pub above)

alk = [HCO3^-(aq)] +2 [CO3^2-(aq)] + [B(OH)₄^-(aq)] + other minor bases

and not as just [OH(aq)^-], [OH(aq)^-] being just another minor base for ocean geochemists.