This post won't say much about the actual issue of ocean acidification. Rather it's a response to something that seems to always happen when the topic arises. Someone inevitably contends that, no, the ocean can't be acidifying because it's alkaline. pH>7 and all that.
There was a long discussion here, for example. What prompted me to post was seeing it happen recently on Judith Curry's blog. It's a sterile topic - a bit like arguing about whether the greenhouse effect is well named. But there's a bit of science in it. The argument can be answered on four levels.
The practical argumentThe meaning is well-known. CO2 in solution will tend to dissolve calcium carbonate, thus disrupting life-forms. It may have other biological effects.
As with the greenhouse effect, human discourse doesn't require literal exactness. We can speak of currency inflation without arguing about whether dollar bills are getting bigger. etc.
The language argumentThe objection is that you can't acidify something if it doesn't become "acid". But that just isn't normal usage. If you beautify something, it doesn't have to become beautiful. We can be enriched without becoming rich.
The theoretical chemistry argumentOK, getting more substantive. The notion that acidic is identified with pH<7 invokes an old notion of acidity. Since about 1923, the more general process going on in acid-base chemistry has been recognised as sharing an electron pair, rather than anything specific with protons. When sulphur trioxide reacts with calcium oxide to produce calcium sulphate, it is easy to recognise this as an acid-base reaction, with SO3 as the acid. No hydrogen is involved.
This is pretty much the case in the ocean. The overall reaction is something like:
CO2 + CaCO3 + H2O → Ca++ + 2HCO3-
Water is a reagent, and there may be a role for protons. But it isn't clear why pH 7 should matter in any way.
The aqueous chemistry argument - bufferingpH 7 is the neutral point of a particular acid-base equilibrium - in pure water. It also applies when a strong acid neutralises a strong base. But the ocean is not pure water, and does not have strong acids or bases.
A solution with substantial concentration of a weak acid and its corresponding base is described as buffered. That is because the pH is stabilised near the neutral point (pKa) of that equilibrium. If you add a strong acid, the protons will react with the corresponding base to produce more of the weak acid. Similar if you add a strong base. The pH changes little, even though real acid-base reaction occurs.
That's why pH 7 is irrelevant here. The ocean is dominated by a 3-way buffering involving CO2, bicarbonate and carbonate ions. There is a further solubility equilibrium between Ca++, CO3-- and solid calcium carbonate (eg aragonite). Adding CO2 pushes everything in the acid direction, which reduces CO3-- concentration and tends to dissolve CaCO3. It's a bit more complicated because the solubility equilibrium is not exactly attained - it is fairly easy for CaCO3 to be supersaturated in solution. But that is the direction.