Recently, Calera proposed a process suitable for power plants and other major CO2 emitters which converts CO2 to calcium carbonate (CaCO3). From the Scientific American:
"
...The Calera process essentially mimics marine cement, which is produced by coral when making their shells and reefs, taking the calcium and magnesium in seawater and using it to form carbonates at normal temperatures and pressures. "We are turning CO2 into carbonic acid and then making carbonate," Constantz says. "All we need is water and pollution...
....Calera hopes to get over that hurdle quickly by first offering a blend of its carbon-storing cement and Portland cement, which would not initially store any extra greenhouse gases but would at least balance out the emissions from making the traditional mortar. "It's just a little better than carbon neutral," notes Constantz, who will make his case to the industry at large at the World of Concrete trade fair in February. "That alone is a huge step forward.".."
In the following, I discuss some of the critical challenges to make this a feasible operation. The roles of the
abundance of calcium and magnesium in sea water and its
pH are discussed in detail. The formation of calcium/magnesium carbonates from sea water
requires energy to proceed. This energy could be supplied either in the form of a
pH shift (by adding strong base) or by
coupling the carbonate formation to other processes (algal photosynthesis).
Why does blending calcium carbonate in cement make economic sense? The unit price for medium ground (4-9 um) CaCO3 is around 95-100 $/T whereas portland cement retails at ~180 $/T . Each ton of CaCO3 blended will save 80 $.Sea water has 0.01 gram-mol Ca2+/kg-sea water and 0.05 gram-mol Mg2+/kg-sea water . The precipitation chemistry of calcium and magnesium carbonates is complicated, involving many minerals. If we precipitate all of the Ca2+ and Mg2+ as carbonates (CaCO3 and MgCO3), sequestering 1 T of CO2 would require 360 T of sea water. Additionally, carbonate precipitation does not occur below a pH of ~10, whereas sea water has a pH of ~7.5-8.4 , which could be decreased by increasing the partial pressure of CO2. Increasing the pH of the solution to favor carbonate precipitation likely requires the use of a base, such as sodium hydroxide, NaOH.Apart from the large volumes of sea water to be handled to make carbonates, the manufacture of NaOH involves the electrolysis of brine and is fairly energy-intensive, producing 4.6 T CO2/T NaOH. Depending on the extent of NaOH requirements, this could represent a significant source of CO2 emissions. (Therefore, using NaOH to sequester CO2 does not make much sense).Switching gears to learning from nature, certain coralline algae form CaCO3 from sea water. This CaCO3 acts as a binder, stabilizing the coral reef. The mechanisms involved are fairly complex for my understanding, but an essential process is the equilibrium between the bicarbonate and carbonate ions, depending on the pH.
pH = 10.33 - log([HCO3-]/[CO32-])Under sunlight, algae photosynthesize dissolved CO2 or HCO3- to form organic compounds (CH2O). This leads to an increase in the local pH. An example of how this happens is given in a paper by John Dodson (subscription required). This increased pH combined with the presence of ion-selective membranes, which admit only certain ions might be how algae use photosynthesis to drive an endothermic CaCO3 precipitation reaction.
Conclusions:- Large volumes of sea water are required to sequester CO2 as calcium/magnesium carbonates. Membrane separations will likely play a critical role in reducing this requirement.
- On the other hand, CO2 sequestration in brines (by product of oil and gas production) also has been investigated. The advantage here is that these brines have high concentrations of calcium, iron and magnesium ions. Therefore, potentially less volumes of brine are needed compared to sea water to sequester the same amount of CO2. The obvious problem here is to increase the pH of the solution to precipitate dissolved CO2 as carbonate. Instead of using a strong base such as sodium hydroxide, folks at NETL used residues from bauxite (an ore of aluminium) processing to buffer the solution pH.
- One needs a driver to precipitate calcium/magnesium carbonates from sea water. In the chemical process described above, it is the addition of sodium hydroxide.
- The smarter way nature has found around that requirement is the formation of carbonates via the bicarbonate<->carbonate equilibrium, shifted towards the right by algal photosynthesis (a reaction converting bicarbonate/CO2 to organic carbon using sunlight, thereby increasing the pH).
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