Mostly Harmless Science Blog: Royal Society Discussion Paper, Ocean acidification due to increasing atmospheric carbon dioxide. Part One.

My thoughts keep returning to the ‘de-alkalinisation of the oceans’. I started thinking about this the other day, first because I came across that article on coccolithophores in Science, and second because one of my students is writing a review article on the use of polymer additives to stop scale formation in desalination plants. The main scales formed in these plants are calcium sulfate at high temperatures, but at somewhat lower temperatures calcium carbonate or magnesium hydroxide.

The first thing you want to know about, if you want to stop scale forming, is what are the characteristics of the solution it is forming from. So early on in the draft appears this table:

(TDS is ‘total dissolved solids’.)

I went back and had another look at the Royal Society discussion paper that I referenced before. This is the paper referenced everywhere in the web where people are fretting about ocean de-alkalinisation. The range of pH values quoted in this table is greater than the range shown in the pretty map in the Royal Society report. In fact, the range of pH values in this table is greater than the size of the maximum change in surface water pH they predict for Figure 5.

So my first thought was, if changes in surface seawater alkalinity are likely to cause bad effects, we ought to be able to see these effects already in ‘canary in the coalmine’ water bodies- shallow, warm places like the Persian Gulf. The reefs there don’t seem to be in particularly good shape but there doesn’t seem to be any evidence that seawater alkalinisation is contributing to their woes. Anyway, this table got me thinking about the problem again.

In discussing the formation of calcium carbonate scale, my student had to talk about the dependence of the equilibrium constants K₁ and K₂ on temperature and the total ionic strength of the solution, and had referenced this paper by Millero et al., where the following figure appears:

The Millero et al. paper also summarises data from a lot of previous work and gets it all to fall on the same line- see this, for instance:

In case you don’t remember,

pK_a = –log₁₀(K_a),

and in this case, K₁ is the equilibrium constant for the reaction:

H₂CO₃ → HCO₃^– + H⁺

and K₂ is the equilibrium constant for this reaction:

HCO₃^– → CO₃^2– + H⁺

These figures are telling us that in seawater (where I^0.5 ~ 0.83), the equilibrium position of both these reactions is further over to the right hand side than if they were happening in common or garden distilled water. And they also tell us that the warmer the water, the further the equilibrium will be over to the right hand side as well.

I plotted up a graph showing how the speciation of pH should change in seawater using the values in this paper and got this figure:

The Royal Society Figure 2 is pretty much the same as mine. It shows carbonate kicking in at a slightly lower pH, but there are different K₂values floating around in the literature and I'm not sure what value they used.