If you’ve ever seen the White Cliffs of Dover you’ll know they’re quite a sight, at least I would imagine. I’ve only seen pictures but even then they’re something. These white monoliths plunge out of the sea reaching heights in excess of 100 m and stretch some 16 km. Perhaps the most striking feature, their stark white colour, is due to the high chalk or limestone composition of the cliffs.
If you took a 100 m3 slice of the cliffs, within that slice you would have something on the order of 2 million tons of CO2 [1]. It’s often we think of CO2 as a problem, but if you look at those cliffs and imagine all that locked up CO2 it makes one think. Since we use limestone, chemically known as Calcium Carbonate, as a product it begs the question, what if CO2 wasn’t a problem? What if it was the solution?
For a moment let’s flip the equation on CO2 and imagine it not as a by-product but as a feed-stock. To do this requires a bit of reverse engineering, that is we’re going to start with the solution and work backwards. In this case the solution is Calcium Carbonate, though it could just as easily be Sodium Carbonate or some other form of Carbonate. These Carbonates are useful in industrial processes; Sodium Carbonate is used in the making of glass, Sodium Bicarbonate is baking soda (nuff said) and Calcium Carbonate is a component or filler in paper, plastics, sheet rock, cement and paints to name a few.
So if our goal is to make Calcium Carbonate how do we get there? Well, you could mine it, preferably not from the Cliffs of Dover. But for the sake of argument if we wanted to make it from its more basic components let’s look at what that would require. We obviously need some Calcium but the other part is Carbonate which you may be a little less familiar with. Looking at the chemical formula for Calcium Carbonate we see it’s CaCO3 so the other part of that equation, the carbonate, is CO3 which looks remarkably similar to CO2, this is no coincidence. When you bubble CO2 through water some will dissolve into the water forming Carbonic Acid (H2CO3). If we increase the pH (i.e. add some Hydroxide ions OH-) we steal some of the Hydrogen ions from the dissolved Carbonic Acid and create Bicarbonate (HCO3-) or if we steal enough Hydrogens we make Carbonate (CO3^2-). So under the right conditions CO2 becomes CO3 and we have the other half of our equation. This is convenient because as we’re all too aware CO2 is a by-product of combustion, and we combust a lot so we have a lot of CO2 for our process, we just need to capture it. Luckily there are already processes that absorb CO2, usually into an Amine stream (Amine is similar to house hold ammonia). Turns out Amines have those Hydroxide ions I was just talking about floating around so it all just ends up being super convenient, turns out if we bubble CO2 through an aqueous basic solution like an Amine we capture CO2 as Carbonate ions.
So I’ve mentioned we need Calcium but I haven’t discussed where we might find it. Calcium is about 4% by mass of the earth’s crust, I’m not going to do the math here just trust me there’s a lot of Calcium out there. Much of that Calcium is already locked up as Calcium Carbonate (limestone), if we heat this limestone up or acidify it we just end up releasing CO2 along with the Calcium so for our purposes that Calcium will be considered off limits. However, anyone who lives with hard water and finds themselves scraping off mineral deposits from their taps and sinks is familiar with the fact Calcium can be found in other places other than locked up in limestone. I’ll dig into some further details shortly but for now let’s presume we can get Calcium from somewhere. With that in mind our process goes something like this; bubble CO2 through a basic solution and create Carbonate ions then dump some Calcium in the mix. The Calcium and Carbonate will react forming Calcium Carbonate which is only slightly soluble in water so most of it precipitates out of solution as a white powder. So in theory we’ve now got a saleable product, Calcium Carbonate in this case but the rational is similar if you chose to go with Sodium Carbonate (or even Magnesium Carbonate). In this case CO2 is not waste, but feed, it’s not the problem, it’s the solution.
This is of course all very neat and tidy but as always the devil is in the details. Even if you have clean sources of Calcium and Carbonate ions you need to combine them in a clever way so you don’t create wasteful by-products. You also have to ensure you don’t require large amounts of another chemical to make things happen since this would be costly. These happen to be just some of the problems that contenders in COSIA’s XPrize are dedicated to solving [2]. I’ve really only skimmed the surface thus far but just for fun let’s take this further and conduct some back of the envelope calculations to see what might be possible. Now be warned, fun is a relative term here, some numbers lay ahead.
First we need to know how much Calcium we need to feed our process. To do this we will have to look at some chemistry, so let’s assume I can remember my high school science and that the general formation of CO3^2- ions goes something like this:
CO2 + H2O -> H2CO3 (some CO2 dissolves in an aqueous solution)
H2CO3 -> 2H+ + CO3^2- (in the presence of a base)
Calcium ions then react with carbonate ions as so:
Ca2+ + CO3^2- -> CaCO3
This has been simplified a bit, in reality this is a multi-step dissociation but for brevity and clarity I think it still works. Using the simplified equations our end result still indicates that for every mol of CO2 consumed we can generate 1 mol of Calcium Carbonate. Now most non-chemical engineers deal in mass not moles so we’ll need to convert, but first we need to determine how much CO2 we have at our disposal to feed the process. Since I’m in Alberta I’ll take a look at a SAGD oil facility. If we’re talking decent reservoirs and a facility with co-generation we can take a CO2 emission intensity of 60 kg of CO2 emitted for every barrel of oil [3]. As a base facility size I’ll assume 10,000 bbl/d so this would produce 600 t/d of CO2 emissions. If you want to lock that up with Calcium take the molecular weight of CO2 at 44.01 g/mol, the molecular weight of CaCl2 at 110.98 g/mol and we know that it’s a 1:1 molar ratio so you need 110.98/44.01 ~ 2.5 times the amount of Calcium Chloride by mass to lock up a kg of CO2. If you wanted to mineralize all the CO2 from a 10,000 bbl/d facility you’re looking at about 1500 t/d of Calcium Chloride. No small amount.
You may have noticed I pulled a bit of a fast one there switching from Calcium to Calcium Chloride, I did so because Calcium doesn’t really exist on its own. It’s too reactive on its own so its always locked up with something. There are mineral sources of Calcium but as mentioned we can’t use limestone, we need silicate rocks like granite or wolllastonite that have Calcium bound to non-carbonate species. Calcium from these mineral sources commonly finds its way into rivers, oceans and aquifers as a salt, CaCl2 (the salt is commonly occurring which is why I used this in the calculation). To get this Calcium you could solution mine deep saline aquifers, there’s apparently 12.9 cubic km’s of saline ground waters in the US [4]. Not all of these will be of a useful concentration but it’s not difficult to find sources with 80,000 mg/l CaCl2 [5]. Taking this concentration you would need to process around 19,000 m3/d of water. That is a lot so you would likely need to process this by some solar evaporation means, using a thermal evaporation process will blow your carbon savings out the window.
Seawater is another source of Calcium Chloride, it’s typically less concentrated at around 500 mg/l so you need to process a lot more water but there is already an industry that desalinates seawater, something like 77 million cubic meters per day in 2011 [6] so the waste streams here could be used. It’s also worth bearing in mind that we’ve just been looking at Calcium salts, if you also consider Sodium and Magnesium salts there’s potential value in those Carbonate forms as well and the amount of water you need to process reduces significantly.
There’s still a lot of work and research to be done to get a clearer picture as to the costs. An old Alberta study from 1993 looked at whether it was feasible to make money solution mining deep saline aquifers for minerals [7]. Turns out it wasn’t. That study looked at CaCl2 and MgCl2 and assumed thermal evaporation methods (i.e. expensive and energy intensive). Our end product here is CaCO3 which depending on the grade may have no additional value or can be twice the value of CaCl2. Precipitated Calcium Carbonate (PCC) is typically a higher grade product as it allows for more control of crystal growth and size. The method I’ve been describing is a precipitation process so there is potential to leverage this fact to obtain a higher end product. Not only has technology changed since 1992 but with a price on carbon and a cap on carbon emissions in industries like the oil sands I’m sure the economics would look different now.
Looking at CO2 as the answer rather than the problem is a simple change in perspective with potentially huge implications. I’ve taken a very high level look here but it would be interesting to see what the energy balance and economics on this looks like but perhaps that’s another blog.
Written by Brent Scarff, P.Eng.
Originally published on LinkedIn on May 23rd, 2017
References:
1. Witty, M. The White Cliffs of Dover are an Example of Natural Carbon Sequestration Ecologia [Online], 2017, p. 23-30. http://scialert.net/fulltext/?doi=ecologia.2011.23.30&org=10.
2. Carbon XPRIZE. http://www.cosia.ca/carbon-xprize.
3. McWhinney, R. Oil Sands Environmental Impacts; Canadian Energy Research Institute: 2014.
4. Howard Perlman, U. Groundwater depletion, USGS water science. https://water.usgs.gov/edu/gwdepletion.html.
5. Hitchon, B.; Bachu, S.; Underschultz, J. R.; Yuan, L. P., Industrial Mineral Potential of Alberta Formation Waters. 1995.
6. Extraction of Formation Water From CO2 Storage. IEAGHG: November 2012.
7. Cross, D. B., Economic Analysis of Extracting Calcium Chloride and Magnesium Chloride from Alberta Brines. March 1993.
