By Dallas Weaver, Ph.D.*
Almost a decade ago, oyster spat production in aquaculture hatcheries in the Pacific Northwest crashed. There was panic in the industry, as even the natural set of baby oysters crashed. Of course, when larval animals die, you can probably find some bacteria associated with the dead animals, however, you seldom know for a fact whether that bacteria was the sole cause or whether it was an opportunist who joined the party which consisted of a bunch of very stressed larvae. As most of us would have done, the oyster culturists chased that bacterial lead into the ground before they examined whether the ocean intake water was changing.
The oyster industry had noted the pH of the water was lower and did some initial experiments to correct the pH, but that didn’t fix the problem. Even with the pH and DO being fixed (which would increase the carbonate concentration), the water didn’t work. However the same water flowing over eel grass adding O2 and removing CO2 did produce good water in the afternoon that would get the larva through the initial stages of fast shell growth.
The solution to this mystery required a lot of independent observations, measurements and concepts that ultimately all tie back to carbonate chemistry. To present this story and its resolution, we need various bits of diverse information, all directly or indirectly related to carbonate chemistry in water.
At the same time as the spat failures, articles appeared in Science (Chan, Barth et al. 2008) about how winds have shifted and created stronger upwellings of deep low oxygen water along the west coast of the US into shallow coastal waters. The low DO of deep water is a product of CO2 being added to the water by the oxidizing of organic materials raining down from algal blooms on the surface. This removes the O2 and increases CO2 in the deep water. When the wind changes, mixing occurs and low DO water comes to the surface near shore. This made its way into the bays, where the oyster hatcheries and production were located. The hatcheries used aeration so the low DO was not an issue.
However, along with low O2 were corresponding higher CO2 levels, which resulted in low pH values. This phenomenon shoved the carbonate/bicarbonate towards lower carbonate. Effectively carbonic acid was being added to the system, which decreased the pH and shifted CO3-- + H+ O3 HCO3- to bicarbonate ion. They were detecting intake pH levels as low as 7.6.
Forms and solubility of limestone:
CaCO3 can exist in many forms, with calcite being the most stable and least soluble. Next in line is aragonite, which is slightly more soluble and amorphous forms that are even more soluble. Animals like oysters often use aragonite in shell building. The equilibrium solubility of aragonite can be represented by:
1) [Ca]* [CO3] = Kar
where: [Ca] and [CO3] are the concentrations of calcium and carbonate and Kar is the apparent solubility constant for aragonite.
In real water systems, activity coefficients, ionic strength, etc. complicate things, but the basic concept remains the same. We can explicitly include these activity coefficients on the concentration terms with their dependence on the other ions in the solution, or they can be combined into an “apparent” solubility product K value.
Nucleation and crystal growth:
In addition, lots of sparingly soluble materials, such as quartz and limestone, only slowly reach chemical equilibrium. When a crystal grows, atoms/molecules are added to a seed crystal and fit into the crystal structure at an atomic level, but when there are no seed crystals or other materials which can match the required atomic spacing and bonding, thermodynamic equilibrium won’t be achieved. However, when the amount of excess calcium and carbonate in solution becomes large enough ([Ca]*[CO3] > > Kar ), the probability of homogenous formation of very small seed crystals increases. Then the solution approaches equilibrium very rapidly with lots of seed crystals. When you mix a strong base, like NaOH or Ca(OH)2 or Na2CO3, with seawater, you can often see the water go white due to the presence of very small crystals being created.
In figure 1, I sampled the water from a small pond where I increased the Ca concentration (with CaCl2) a month ago to 550 mg/l at a pH of 8.24 with an equilibrium pH of 8.45. The pond had a small amount of additional pCO2 from decomposing leaves. Both samples were the same water, but the water in the cloudy sample contained the addition of 3 % of a 1 % Na2CO3 solution, this increased the pH from 8.24 to about 8.5. However, the local area where the solutions mixed created a precipitate that worked as a nucleus for the precipitation of CaCO3 from solution. On aerating both samples for about 2 hours, in the cloudy sample, to which I added alkalinity as washing soda, the pH decreased to 8.13 and the measured Ca decreased to 350 (with more time, it will continue to rapidly decrease, but the sample without nucleation won’t change in months). The untreated sample increased in pH to 8.45. When we are dealing with non-equilibrium chemistry and kinetics, the devil is in the details.
What I did was reduce the [Ca] concentration and the alkalinity by adding alkalinity. This may sound paradoxical, but [Ca] and [Mg] (hardness) are commercially removed from water by adding calcium as Ca(OH)2 (hydrated lime) and Na2CO3 (washing soda). This shifts the carbonate chemistry towards [CO3] and [OH] while decreasing the solubility of the [Ca] and creating seed crystals for the CaCO3 to precipitate on. This is referred to as lime/soda softening and is a very clever carbonate chemistry manipulation game.
The odd behavior of limestone and supersaturation:
Limestone also has the property of being more soluble under greater pressure and lower temperature (unlike most materials, which are more soluble at higher temperatures). That means that limestone in the deep ocean (high pressure, low temperature) will dissolve. Then, when global ocean circulation finally returns the water back to the surface, where it warms up, the water can be highly supersaturated, containing several times more calcium than would be in equilibrium with the amount of carbonate present in the oceans. This phenomenon creates carbonate scale in hot water pipes in large of areas of the country, which in some areas is very useful in protecting the water from absorbing lead contained in old water pipes.
We can define that supersaturation as Ωar for the ratio of the actual [Ca]*[CO3] / Kar. In one sense, supersaturation (Ω > 1) is a form of chemical energy that animals, such as corals and larval oysters, can use to build their shells or skeletons without spending energy pumping ions. Under these non-equilibrium conditions, a crystal will naturally grow if a crystal seed is present, so all the animal has to do is produce a small amount of organic material which matches the desired crystal structure well enough for the crystal growth to start. In essence, the animal creates an artificial organic “seed” crystal. The animal gets its shell for a very small energy investment. When 80 % of a larval oyster’s body weight is CaCO3, it can’t afford to pump ions around.
Dallas Weaver, PhD, started designing and building closed aquaculture systems in 1973 and worked for several engineering/consulting companies in the fields of air pollution, liquid wastes, and solid wastes until 1980. Today, he’s the Owner/President of Scientific Hatcheries.